INTRODUCTION  TO 
GENERAL  CHEMISTRY 


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INTRODUCTION 

TO 

GENERAL  CHEMISTRY 


BY 

HERBERT  N.  McCOY 


ETHEL  M.  TERRY 


SECOND  EDITION 
FIRST  IMPRESSION 


McGRAW-HILL  BOOK  COMPANY,  INC. 

NEW  YORK:    239   WEST  39TH  STREET 

LONDON:    6  &  8  BOUVERIE  ST.,  E.  C.  4 

1920 


0\ 


COPYRIGHT,  1919,  1920,  BY 
HERBERT  N.  McCov  AND  ETHEL  M.  TERRY 


PREFACE 

This  text  has  been  in  process  of  growth  at  the  University 
of  Chicago  since  1913.  At  that  time  a  synopsis  of  the  first 
nine  chapters  was  printed.  This  was  followed,  in  1 9 1 6-1 7 ,  by  the 
completed  work  of  the  first  fifteen  chapters.  Chapters  xvi-xix 
were  added  and  used  in  class  work  during  1918  and  1919. 

The  book  has  been  written  for  college  Freshmen,  and,  as  its 
title  implies,  it  is  intended  to  serve  as  an  introduction  to  general 
chemistry.  In  consequence  we  have  aimed  to  present  a  contin- 
uous and  connected  story  in  teachable  form  and  have  not 
attempted  to  give  extensive  descriptive  and  numerical  data 
where  such  matter  is  of  little  interest  to  the  student  or  is  not 
needed  for  the  development  of  important  principles. 

Inasmuch  as  the  choice  and  arrangement  of  topics  in  the 
earlier  part  of  this  book  depart  noticeably  from  the  familiar 
order,  some  explanation  seems  necessary.  We  shall  therefore 
sketch  briefly  the  plan  of  the  less  conventional  chapters, 
together  with  the  " philosophy  of  arrangement"  which  has 
resulted  in  the  scheme  presented. 

The  first  chapter,  which  is  brief,  deals  with  the  measurement 
of  gases  and  the  gas  laws.  In  the  next  five  chapters  the  most 
fundamental  concepts  of  the  science  of  chemistry  are  developed. 
These  include:  indestructibility  of  matter,  idea  of  a  pure  sub- 
stance, decomposition  of  pure  substances,  elements,  analysis  of 
substances  and  percentage  composition,  the  law  of  definite 
composition,  derivation  of  formulae. 

Chapter  v  shows  how  chemical  formulae  are  derived  from  a 
knowledge  of  percentage  composition  and  gas  or  vapor  density. 
This  development  keeps  as  close  as  possible  to  the  arguments 
of  Avogadro  and  Cannizzaro  and  shows  how  formulae  are  ob- 
tained by  methods  independent  of  the  atomic-molecular  hypoth- 
esis. Reference,  at  this  stage,  to  combining  weights  and 
chemical  equivalents  is  purposely  avoided,  for  the  reason  that 
the  history  of  chemistry  between  the  time  of  Avogadro  and 


415329 


vi  Preface 

that  of  the  epoch-making  paper  of  Cannizzaro  (1858)  shows 
the  fallacy  of  trying  to  develop  formulae  and  fix  atomic 
(symbol)  weights  by  any  method  other  than  that  proposed  by 
Avogadro  and  elucidated  by  Cannizzaro.  The  sixth  chapter 
introduces  the  use  of  equations  before  the  atomic-molecular 
hypothesis  is  studied.  This  plan  has  the  great  advantage  of 
fixing  in  the  student's  mind  the  fundamental  relationship 
between  equations  and  the  quantitative  experimental  data 
such  equations  represent. 

Chapters  ii  to  vi  inclusive  form  a  compact  division  of  the 
subject,  in  which  the  argument,  illustrated  at  every  step  by 
experimental  data,  is  substantially  continuous.  In  these  chap- 
ters we  have  aimed  at  a  logical  development  of  the  subject 
without  the  introduction  of  any  matter  that  does  not  serve  to 
illustrate  the  topics  under  discussion. 

The  next  three  chapters,  vii,  " Acids,  Bases,  and  Salts — I"; 
viii,  "Water  and  Solutions";  and  ix,  " Acids,  Bases,  and  Salts— 
II,"  are  introduced  at  this  point  for  very  definite  reasons.  In 
the  first  place,  it  is  obvious  to  every  teacher  that  much  of  the 
beginner's  work  will  deal  with  acids,  bases,  and  salts  and  their 
solutions.  It  is  our  opinion  that  a  knowledge  of  these  topics 
is  best  obtained  by  studying  them  directly  and  specifically, 
both  in  the  classroom  and  in  the  laboratory.  Our  plan  provides 
for  laboratory  work  by  the  student,  following  closely  the  content 
of  these  three  chapters.  This  laboratory  work  is  interesting  to 
students,  since  they  like  to  make  and  crystallize  a  variety  of 
salts.  It  also  gives  good  training  in  technique  and  is  not 
difficult  either  experimentally  or  theoretically,  while  at  the 
same  time  it  offers  a  wealth  of  material  for  practice  in  writing 
equations  and  solving  problems.  The  most  important  reason, 
however,  for  the  introduction  of  the  early  study  of  acids, 
bases,  and  salts  is  to  supply  the  indispensable  data  needed  later 
for  the  understanding  of  the  ionic  hypothesis. 

Chapters  x  and  xi  present  the  kinetic-molecular  and  atomic 
hypotheses  respectively.  It  will  be  noted  that  these  subjects 
follow  the  development  and  use  of  formulae  instead  of  preceding 
them.  This  emphasizes  the  generally  overlooked  fact  that  for- 


Preface  vii 

mulae  are  in  no  way,  of  necessity,  dependent  upon  the  molecular- 
atomic  hypothesis.  At  the  same  time  the  student  is  in  a  position 
to  appreciate  more  readily  the  molecular-atomic  hypothesis  be- 
cause it  furnishes  a  plausible  explanation  of  facts  already  familiar 
to  him. 

Since  the  early  chapters  of  the  text  contain  as  much  infor- 
mation about  hydrogen  and  oxygen  as  the  student  needs,  or 
can  fully  appreciate  at  the  start,  the  formal  discussion  of  these 
elements  has  been  postponed  to  chapter  xiv.  Chapters  xiii 
("Chemical  Equilibrium"),  xv  ("Oxidation  and  Reduction"), 
and  xvi  ("Heat  and  Energy")  present  important  theoretical 
matters. 

The  extended  discussion  given  of  the  ionic  hypothesis  needs 
no  apology.  In  this  connection  we  have  introduced  a  new 
method  of  graphic  representation  of  ionic  equilibrium.  In 
chapter  xx,  on  "Electrochemistry,"  the  electronic  conception 
of  reactions,  including  oxidation  and  reduction,  is  discussed. 

The  arrangement  from  this  point  on  needs  little  comment. 
The  authors  recognize,  of  course,  that  organic  chemistry  (xxv 
and  xxvi),  "Theory  of  Dilute  Solutions"  (xxvii),  "Disperse 
Systems"  (xxviii),  and  "Radioactivity"  (xxxii)  are  optional 
studies  for  beginning  general  chemistry  courses.  We  have 
included  these  subjects,  since  the  interest  which  they  have 
aroused  has  apparently  justified  so  doing. 

The  880  sections  of  the  book  have  been  not  only  titled  but 
numbered,  in  order  that  frequent  cross-references  might  be 
given.  The  student  is  thus  constantly  informed  of  the  mate- 
rial on  which  each  discussion  is  based.  Instructors  will  find  the 
references  helpful  in  laying  out  work  if  they  desire  to  skip 
certain  portions  of  the  text,  for  instance  with  classes  of  students 
who  have  had  a  secondary-school  course  in  chemistry. 

The  great  importance  of  close  connection  between  the  work 
of  the  laboratory  and  that  in  class  has  been  kept  constantly 
in  mind  in  planning  this  text.  A  laboratory  guide  (4  Labora- 
tory Outline  for  General  Chemistry,  by  McCoy  and  Terry),  which 
follows  strictly  the  arrangement  of  the  text,  has  been  written 
to  accompany  the  latter. 


viii  Preface 

In  addition  a  pamphlet  has  been  prepared  containing  those 
details  of  the  lecture  experiments  which  are  important  to  the 
lecture  assistant  but  of  no  interest  to  the  elementary  student. 
This  pamphlet  will  be  available  for  teachers  using  the  text. 

We  are  indebted  to  Professor  W.  D.  Harkins,  of  the  Uni- 
versity of  Chicago,  for  the  contribution  of  sections  6  and  7 
(pp.  3  and  4)  of  the  text  and  for  the  use  of  Fig.  117;  to  Mr. 
Leo  Finkelstein  for  the  drawings  of  Figs,  i  to  43;  and  to  Dr. 
R.  D.  Mullenix  for  the  seventy-seven  additional  drawings,  as 
well  as  for  valuable  criticism. 

THE  AUTHORS. 

CHICAGO,  ILL. 
July,  1919 


NOTE  TO  SECOND  IMPRESSION. 

We  have  taken  advantage  of  the  necessity  of  a  second 
printing  of  this  text  to  correct  a  number  of  misprints  and  to 
add  a  chapter  on  metallurgy.  In  the  preparation  of  the  latter 
we  have  found  the  following  works  of  general  assistance: 
Economic  Geology,  by  Heinrich  Ries  (John  Wiley  &  Sons) ; 
Non-technical  Chats  on  Iron  and  Steel,  by  LaVerne  W.  Spring 
(Frederick  A.  Stokes  Co.);  a  bulletin  entitled  " Copper"  pub- 
lished by  the  Anaconda  Copper  Mining  Co.,  New  York;  The 
Mineral  Industry  in  1918;  Principles,  Operation  and  Products 
of  the  Blast  Furnace,  by  J.  E.  Johnson,  Jr.;  and  the  works 
of  H.  O.  Hoffman  on  General  Metallurgy,  The  Metallurgy  of 
Lead,  and  The  Metallurgy  of  Copper.  (The  last  five  books 
are  all  published  by  the  McGraw-Hill  Book  Co.) 

THE  AUTHORS. 

CHICAGO,  ILL., 
January,  1920. 


TABLE  OF  CONTENTS 

CHAPTER  pAGE 

I.  INTRODUCTION — LAWS  or  GASES i 

II.  THE  BURNING  OF  SUBSTANCES — OXYGEN 6 

III.  PURE  SUBSTANCES — ELEMENTS 16 

IV.  THE  LAW  OF  DEFINITE  COMPOSITION 26 

V.  SYMBOLS  AND  CHEMICAL  FORMULAE 40 

VI.  CHEMICAL  EQUATIONS 5° 

VII.  ACIDS,  BASES,  AND  SALTS — I 57 

VIII.  WATER  AND  SOLUTIONS 68 

IX.  ACIDS,  BASES,  AND  SALTS — II 86 

X.  THE  KINETIC  THEORY  OF  MATTER  AND  THE  MOLECULAR 

HYPOTHESIS 109 

XI.  THE  ATOMIC  HYPOTHESIS  AND  ATOMIC  WEIGHTS  .     .     .  121 

XII.  THE  HALOGENS  AND  THEIR  COMPOUNDS  WITH  HYDROGEN 

AND  METALS 135 

XIII.  CHEMICAL  EQUILIBRIUM 158 

XIV.  HYDROGEN  AND  OXYGEN 176 

XV.  OXIDATION  AND  REDUCTION 194 

XVI.  HEAT  AND  ENERGY 214 

XVII.  THE  IONIC  HYPOTHESIS 227 

XVIII.  APPLICATIONS  OF  THE  IONIC  HYPOTHESIS 251 

XIX.  APPLICATIONS    OF  THE  IONIC   HYPOTHESIS.    REACTIONS 

INVOLVING  CHANGES  OF  STATE 275 

XX.  ELECTROCHEMISTRY  . 296 

XXI.  NITROGEN  AND  AMMONIA 323 

XXII.  NITRIC  ACID  AND  THE  OXIDES  OF  NITROGEN  .     .     .     .341 

XXIII.  PHOSPHORUS  AND  ITS  COMPOUNDS 363 

XXIV.  SULFUR  AND  ITS  COMPOUNDS 375 

XXV.  CARBON  AND  CARBON  COMPOUNDS,  ORGANIC  COMPOUNDS — 

I 393 

ix 


x  Table  of  Contents 

CHAPTER  PAGE 

XXVI.  ORGANIC  COMPOUNDS — II 426 

XXVII.  THEORY  OF  DILUTE  SOLUTIONS 452 

XXVIII.  DISPERSE  SYSTEMS 470 

XXIX.  THE  ATMOSPHERE  AND  RELATED  TOPICS 491 

XXX.  SOME  ADDITIONAL  ELEMENTS  AND  THEIR  COMPOUNDS     .  514 

XXXI.  CLASSIFICATION  OF  THE  ELEMENTS.    THE  PERIODIC  SYSTEM  539 

XXXII.  RADIOACTIVITY  AND  THE  NATURE  OF  MATTER.      .     .     .  568 

XXXIII.  METALLURGY       . 59* 

INDEX                                                         633 


CHAPTER  I 


INTRODUCTION— LAWS  OF  GASES 

1.  A  Knowledge  of  Physics  Prerequisite  for  Chemistry.-— 

The  sciences  of  physics  and  chemistry  are  so  closely  related  that 
the  latter  may  be  considered  an  extension  of  the  former.  A 
knowledge  of  physics  is  therefore  necessary  for  an  adequate 
understanding  of  chemistry,  and  it  is  to  be  assumed  that  the 
student  taking  up  chemistry  has  had  at  least  a  one-year  high- 
school  course  in  physics. 

2.  The  Three  Forms  of  Matter:    Gases. — In  his  work  in 
physics,  the  student  will  have  learned  the  meaning  of  the  term 
matter,  which  may  be  denned  as  anything  which  occupies  space 
and  has  weight.    He  will  have  learned,  also,  that  matter  may 
exist  in  three  forms:  solid,  liquid,  and  gaseous.     Since  gases  are 
less  tangible  than  solids  and  liquids,  we  shall  first  take  up  the 
study  of  air,  the  most  familiar  of  all  gases.     That 

air  has  the  two  attributes  just  mentioned  as  belong- 
ing to  all  forms  of  matter  may  readily  be  shown  by 
experiment. 

3.  Air  Occupies  Space  and  Has  Weight. — If  a 
drinking  glass  or  beaker  be  thrust,  mouth  down- 
ward, into  a  vessel  of  water,  the  water  does  not 
enter  until  the  glass  is  tilted  to  allow  the  air  to 
escape.     This  shows  that  air  occupies  space. 

That  air  has  weight  may  be  shown  by  weighing 
a  flask,  first  empty  and  afterward  filled  with  air. 
The  flask  (Fig.  i)  should  be  round-bottomed  and 
have  a  capacity  of  250  to  500  c.c.  It  is  fitted  with 
a  tight  rubber  stopper  carrying  a  glass  stopcock.  The  air  is 
first  pumped  out  by  means  of  an  efficient  air  pump;  the  stop- 
cock is  then  closed  and  the  flask  counterbalanced  with  weights. 
When  the  stopcock  is  opened  the  inrush  of  air  can  be  heard, 
and  it  is  easy  to  observe  that  there  is  an  appreciable  increase 


FIG.  i 


2  Introduction  to  General  Chemistry 

in  weight.     Since  air  occupies  space  and  has  weight,  it  is  un- 
doubtedly a  form  of  matter. 

One  liter  of  air  weighs  more  than  a  gram  and  the  air  contained 
in  a  room  12  feet  square  and  12  feet  high  would  weigh  about  100 
pounds.  At  the  earth's  surface  air  exerts  a  pressure  of  about 
15  pounds  on  every  square  inch  of  surface.  The  existence  of 
this  pressure  may  readily  be  shown  by  means  of  the  following 
experiment.  A  tin  can  with  a  narrow  neck  (such  as  is  often  used 
for  shipping  alcohol,  etc.)  and  of  about  i  gallon  capacity  is 
fitted  with  a  stopper  carrying  a  glass  tube,  by  means  of  which 
the  air  filling  the  can  may  be  pumped  out.  Usually,  before  the 
exhaustion  of  the  air  is  complete,  the  can  is  crushed  by  the 
pressure  of  the  air  on  the  outside — a  pressure  which  is  now 
no  longer  balanced  by  the  equal  and  opposite  pressure  on  the 
inside. 

4.  The  Effect  of  Pressure  on  Volume:   Boyle's  Law. — The 
atmospheric  pressure  is  measured  by  means  of  the  barometer. 
At  the  sea-level  the  normal  barometric  pressure  serves  to  support 
a  column  of  mercury  76  cm.  high.    The  effect  of  pressure  upon 
the  volume  of  air  was  first  studied  by  Robert  Boyle  in  the  seven- 
teenth century.     Boyle  found  that  the  volume  of  a  given  portion 
of  air  was  inversely  proportional  to  the  pressure.     This  relation 
is  known  as  Boyle's  law.     If  we  represent  the  pressure  by  P  and 
the  volume  by  V,  then  PF  =  a  constant. 

5.  The  Effect  of  Temperature   on  Volume:    The  Law  of 
Charles. — In  scientific  work  we  use  the  Centigrade  thermometer, 
the  scale  of  which  is  so  constructed  that  the  freezing-point  of 
water  is  o°,  while  the  boiling-point  is  100°.   The  effect  of  tempera- 
ture upon  the  volume  of  a  given  portion  of  air  at  a  fixed  pressure 
was  studied  over  a  century  ago  by  Charles  and  by  Gay  Lussac. 
It  was  found  that  the  volume  of  the  air  increased  1/273  °f  its 
volume  at  zero  for  each  increase  of  i°  C.     This  statement  is 
known  as  the  law  of  Charles,  or  sometimes  also  as  the  law  of 
Gay  Lussac. 

6.  The    Gas    Thermometer:     Absolute    Temperature. — An 
experiment  will  show  that  if  273  c.c.  of  air  contained  in  a  flask 
or  cylinder  at  o°  C.  is  heated  to  100°  C.  the  volume  will  change  to 


Introduction — Laws  of  Gases  3 

373  c.c.  Such  an  apparatus  is  called  an  air  thermometer  and 
temperatures  may  be  measured  in  this  way  instead  of  by 
the  expansion  of  mercury,  as  in  ordinary  thermometers.  At 
i°  C.  the  volume  of  the  air  is  274  c.c.;  at  2°  it  equals  275  c.c. 
and  thus  the  volumes  in  the  following  table  correspond  to 
the  temperatures  given. 

TABLE  I 


Volume                        Degrees 
in  c.c.                       Centigrade 

Volume                           Degrees 
in  c.c.                          Centigrade 

373  minus  2' 
372 
363       " 
323       " 

73  =  100 
=  99 
=  90 

=   50 

293   mi 
283 
273 
263 

nus  2 

73=     20 

'    =        IO 
'    =          0 
'     =  —  10 

3i3 
303       " 

=  40 
=  30 

253 
243 

<    =—  20 
'    =-30 

Since  the  zero  of  the  Centigrade  thermometer  is  arbitrarily 
chosen,  being  the  temperature  of  the  freezing  of  water  (the  stu- 
dent is  already  familiar  with  the  Fahrenheit  zero,  which  is  at  a 
lower  temperature),  it  would  be  possible  and  convenient  to  use 
a  temperature  scale  in  which  the  volumes  of  the  air  in  the  air 
thermometer  as  described  are  taken  as  the  temperatures.  Since 
temperatures  on  the  Centigrade  scale  are  obtained  by  subtracting 
273  from  the  corresponding  air- thermometer  temperatures,  the 
zero  of  the  air  thermometer  or  gas  scale  must  be  273  degrees  lower 
than  the  Centigrade  zero,  or  273  degrees  below  the  freezing-point. 
These  air-thermometer  temperatures  are  usually  called  the  abso- 
lute temperatures;  the  absolute  temperature  may  therefore  be 
denned  as  the  Centigrade  temperature  plus  273  degrees.  Since 
most  other  gases  act  like  air  they  may  be  used  in  the  gas  ther- 
mometer, and  it  is  evident  that  if  a  certain  amount  of  gas  is  used 
in  such  an  experiment,  no  matter  what  its  volume  may  be  at  the 
freezing-point  of  water,  the  volume  will  always  vary  with  the 
temperature  in  the  same  ratio  as  the  absolute  temperature,  pro- 
vided the  pressure  on  the  gas  is  kept  constant. 

7.  Problems. — We  may  now  consider  a  few  simple  problems 
based  on  the  two  laws  of  gases  just  discussed. 


4  Introduction  to  General  CJiemistry 

Problem  i:  The  volume  of  a  certain  amount  of  air  at  27°  C. 
is  1,000  c.c.  What  would  its  volume  be  at  127°  C.  if  the  pressure 
is  kept  constant? 

Centigrade  temperature+ 2  73=  absolute  temperature 


The  volume  of  the  gas  must  therefore  increase  in  the  ratio  of 
400  to  300,  or  it  will  become 

*  .  400 
loooc.c.X—  =  1333. 3  c.c. 

Problem  2:^  Let  the  original  pressure  on  the  gas  in  Problem  i 
be  60  cm.  of  mercury  (or  —^  of  the  ordinary  pressure  of  the  atmos- 
phere). What  will  be  the  final  volume  of  the  gas  if  the  pressure 
is  increased  to  100  cm.  of  mercury?  An  increase  of  pressure 
must  decrease  the  volume  of  the  gas,  and  in  the  ratio  of  the 

pressures,  60  to  100,  or  by  - 

a)  Let  the  change  of  pressure  come  after  the  change  of 
temperature  as  given  in  Problem  i : 

then 

1333  .3  c.cX  —  =  800  c.c.  (final  volume)    Ans. 

b)  Let  the  change  of  pressure  take  place  first : 

1000  c.cX —  =  600  c.c.,  volume  after  the  pressure  change. 

The  temperature  change  would  then  change  the  volume 'as 
follows: 

6oc  c.c.X- —  =  800  c.c.  (final  volume)     Ans. 
300 

It  is  thus  seen  that  the  same  answer  is  obtained,  no  matter 
which  step  in  the  problem  is  worked  first,  so  the  whole  problem,, 
i  and  2  together,  may  be  stated  in  one  expression  as  follows: 

1000  c.c.  X — X —  =  800  c.c.    Ans. 
300     100 


Introduction — Laws  of  Gases  5 

Problems:  Suppose  that  1,000  c.c.  of  air  at  20°  C.  and  70  cm. 
pressure  is  cooled  to  o°  and  that  at  the  same  time  the  pressure 
is  increased  to  76  cm.  Find  the  final  volume. 

When  a  gas  is  at  the  temperature  of  o°  C.  and  under  a  pressure 
of  76  cm.  (the  normal  atmospheric  pressure  at  sea-level)  it  is  said 
to  be  at  standard  conditions.  Unless  otherwise  stated,  numeri- 
cal data  for  gases  refer  to  the  latter  under  standard  conditions. 

Problem  4:  Find  the  volume  at  standard  conditions  of  400  c.c. 
of  air  measured  at  30°  and  72  cm. 

8.  Steam  Is  the  Gaseous  Form  of  Water. — It  is  well  known 
that  when  water  is  heated  it  passes  into  steam.    The  white  cloud 
which  is  frequently  spoken  of  as  steam  is  not  really  steam,  but 
is  composed  of  minute  droplets  of  water.     If  we  boil  water  in  a 
glass  flask  the  space  above  the  water  is  filled  with  steam,  but 
we  notice  that  the  steam  is  entirely  invisible  and  that  the  visible 
cloud  forms  only  when  the  steam  cools  and  condenses  to  liquid 
droplets.    Water  in  the  form  of  steam  is,  like  air,  a  gas.    When 
we  boil  any  liquid  like  alcohol  or  mercury  the  liquid  passes  into 
the  state  of  a  gas  or  vapor,  as  it  is  sometimes  also  called.    The 
gas  or  vapor  when  cooled  condenses  to  the  liquid  form  of  the  sub- 
stance. 

9.  Change  of  Form  of  Matter  with  Change  of  Temperature.— 
Just  as  water  when  cooled  solidifies  to  ice,  so  every  other  liquid 
substance  solidifies  when  sufficiently  cooled.    We  speak  of  steam 
and  ice  as  the  gaseous  and  solid  forms  respectively  of  water. 
The  substance  known  as  moth-balls  is  called  naphthalene  by  the 
chemist;  it  is  a  solid  at  ordinary  temperatures,  but  when  heated 
it  melts  to  a  colorless  liquid,  and  when  heated  still  hotter  it 
boils,  giving  a  colorless  vapor,  which  is  naphthalene  in  the  form 
of  a  gas.     When  this  gas  is  cooled  it  condenses  to  a  liquid,  which 
when  cooled  still  further  solidifies  or  freezes,  giving  solid  naph- 
thalene again.     Behavior  like  that  of  water  and  naphthalene 
is  met  with  in  the  case  of  very  many  other  substances.     They 
can  exist  in  three  different  forms,  gas,  liquid,  and  solid,  according 
to  the  temperature. 


CHAPTER  II 


i 


FIG.  2 


THE  BURNING  OF  SUBSTANCES— OXYGEN 

10.  Burning  Substances  Require  Air. — The  history  of  chem- 
istry shows  that  the  discovery  of  the  real  nature  of  the  process  of 
burning  was  one  of  the  most  important,  if  not  the  most  important, 
in  the  development  of  the  whole  science.  That  air  is  needed  for 
the  burning  of  a  substance  is,  in  general,  well  known,  and  can 

easily  be-  shown  by  many  simple 
experiments.  For  example,  if  we 
place  an  inverted  drinking-glass 
over  a  burning  candle  standing  on 
a  table  (Fig.  2),  the  flame  quickly 
grows  smaller  and  smaller  and  soon 
goes  out,  the  glass  having  cut  off 
1  the  needed  supply  of  air. 

A  still  more  interesting  and  in- 
structive experiment  may  be  made 

with  phosphorus,  a  substance  which  burns  very  readily  in  the 
air,  giving  off  clouds  of  white 
smoke.  A  piece  of  phosphorus  of 
the  size  of  a  pea  is  placed  on  a 
cork  floating  on  water  and  covered 
with  a  bell-jar  (Fig.  3).  When  a 
heated  wire  passing  through  the 
tight-fitting  stopper  of  the  jar  is 
brought  in  contact  with  the  phos- 
phorus the  latter  takes  fire  and 
burns  with  the  production  of  light 
and  heat  and  the  formation  of  a 
cloud  of  white  smoke.  At  the 
same  time  the  level  of  the  water 

inside  the  bell-jar  first  falls  a  little  and  later  rises;  but  while 
there  is  still  a  large  volume  of  air  left  above  the  water  on  which 
the  cork  floats,  the  flame  dies  out  and  the  burning  ceases.  By 

6 


(  (• 


FIG.  3 


The  Burning  of  Substances — Oxygen  7 

the  time  the  bell-jar  and  its  content  have  become  cold,  the 
cloud  has  disappeared  and  the  water  has  risen  on  the  inside  so 
that  the  volume  of  the  remaining  air  is  seen  to  be  about  four- 
fifths  of  the  original  volume.  It  follows  that  about  one-fifth 
by  volume  of  the  air  has  disappeared. 

Further  examination  also  shows  that  much  of  the  phos- 
phorus still  remains  unburned.  Why,  then,  should  the  burn- 
ing stop  while  there  is  still  four-fifths  by  volume  of  the  air  left 
in  the  jar?  The  answer  to  this  question  may  be  made  when 
we  find  that,  try  as  we  may,  we  cannot  make  phosphorus  or 
anything  else  burn  in  the  air  remaining  in  the  jar.  We  therefore 
conclude  that  the  remaining  air  is  different  from  common  air. 
The  correctness  of  this  conclusion  is  supported  by  the  fact  that 
small  animals,  such  as  mice,  suffocate  at  once  if  allowed  to  breathe 
this  remaining  portion  of  the  air.  The  facts  just  considered  make 
it  seem  probable  that  one-fifth  of  the  air  is  different  from  the 
balance,  and  that  it  is  this  portion  which  takes  part  in  the  burning 
of  substances  and  which  is  necessary  for  the  respiration  of 
animals. 

Everyday  experience  would  seem  to  indicate  that  wood,  coal, 
paper,  gasoline,  etc.,  are  completely  destroyed  when  they  are 
burned.  Wood  and  coal  leave  a  small  amount  of  ash  when 
burned,  but  nothing  visible  remains  in  the  case  of  gasoline  and 
other  oils.  Since  we  have  found  that  water  in  the  form  of  steam 
is  invisible,  it  is  possible  that  the  substance  burned  may  have 
passed  into  an  invisible  form  and  thus  escaped  notice. 

There  are  many  substances  which  burn  very  readily  and  in 
so  doing  leave  behind  large  amounts  of  ash;  the  experimental 
study  of  the  burning  of  such  substances  leads  to  important  con- 
clusions. We  may  now  consider  two  typical  cases  of  this  sort. 

ii.  The  Burning  of  Magnesium. — The  metal  magnesium, 
which  is  used  in  photographic  flash  lights,  will  burn  very  readily 
in  air,  either  in  the  form  of  powder  or  thin  ribbon.  In  either  case 
we  notice  that  a  white  ash  is  left.  If  we  collect  and  weigh  the  ash 
from  the  burning  of  a  weighed  piece  of  magnesium  ribbon  we 
find  that  the  ash  weighs  more  than  the  original  metal  ribbon. 
The  actual  experiment  is  best  carried  out  by  placing  about  one 


8 


Introduction  to  General  Chemistry 


gram  of  magnesium,  in  the  form  of  wire  (Fig.  4)  or  ribbon,  in  a 
porcelain  crucible,  having  a  cover,  and  then  weighing  crucible  and 

contents.  The  magnesium  is  then 
ignited  and  the  cover  so  adjusted 
that  some  air  can  enter,  but  that 
the  dense  cloud  of  white  smoke  is 
largely  held  back  in  the  crucible. 
After  the  burning  is  finished  and 
the  crucible  has  cooled  and  the 
whole  is  again  weighed,  it  will  be 
found  that  there  has  been  a  consider- 
able increase  in  weight. 
12.  The  Burning  of  Iron. — Iron  powder  or  filings  burn  readily 
when  thrown  into  a  flame,  and  in  a  similar  manner  we  find  that 
the  burned  iron  or  iron  ash,  as  we  might  possibly  call  it,  is  heavier 
than  the  original  metal.  In  order  to  show  this  by  experiment, 
we  may  suspend  on  one  side  of  a  balance  (Fig.  5)  a  horseshoe 


FIG.  4 


FIG.  5 


magnet  which  has  been  dipped  in  iron  filings,  and  counterpoise 
the  magnet  and  adhering  iron  by  adding  small  shot  or  sand  to 
the  other  pan  of  the  balance.  By  the  application  of  a  flame,  the 
iron,  which  now  presents  a  large  surface  to  the  air,  may  be 
ignited.  As  it  burns  with  a  dull  glow  we  observe  a  gradual 
increase  in  its  weight,  and,  while  there  is  no  noticeable  change 
in  its  volume,  the  cold  residue,  which  we  may  call  iron  ash,  is 


The  Burning  of  Substances — Oxygen 


seen  to  have  lost  its  metallic  luster  and  taken  on  a  dead  black 
color.  We  find,  thus,  that  iron  ash  is  heavier  than  the  iron  burned. 
If  we  seek  the  cause  of  this  increase  in  weight,  we  may  get  a 
hint  when  we  remember  that  for  the  burning  of  a  candle  air  is 
required,  and  that,  moreover,  part  of  the  air  disappeared  when 
phosphorus  was  burned  in  it.  What,  then,  becomes  of  the 
weight  of  the  one-fifth  of  the  air  that  disappeared?  Is  it  added 
to  the  weight  of  the  iron,  so  as  to  increase  the  weight  of  its  ash? 
The  facts  presented  in  the  next  paragraph  will  furnish  the 
required  answers. 

13.  Lavoisier's  Experiment  with  Mercury0 — An  experiment 
which  turned  out  to  be  one  of  the  most  important  made  in  the 
early  development  of  the 
science  of  chemistry  was 
carried  out  by  the  great 
French  chemist,  Lavoisier, 
in  the  latter  part  of  the 
eighteenth  century.  The 
arrangement  in  this  classic 
experiment  is  shown  in 
Fig.  6.  The  retort  (the 
glass  vessel  with  the  long 
bent  neck)  was  partly  filled 
with  mercury  (quicksil- 
ver) ;  the  space  above  the  mercury  contained  ordinary  air,  which 
also  filled  the  bell-jar  with  which  the  neck  of  the  retort  communi- 
cated. The  bell-jar  stood  in  a  shallow  vessel  containing  mercury, 
which  served  to  prevent  outside  air  from  passing  into  or  out  of 
the  jar.  The  mercury  in  the  retort  was  now  heated  by  means  of 
a  charcoal  stove  for  a  period  of  several  days.  The  heating  first 
caused  an  expansion  of  the  air;  but  as  time  went  on  a  gradual 
contraction  occurred,  which  entirely  ceased  after  several  days, 
whereupon  the  heating  was  stopped.  The  volume  of  the  air  left 
in  the  entire  apparatus  when  brought  to  its  original  temperature 
and  pressure  was  practically  four-fifths  of  what  it  had  been  at  the 
start.  The  surface  of  the  mercury  in  the  retort  was  found 
to  be  covered  with  a  red  powder,  which  may  be  considered 


FIG.  6 


io  Introduction  to  General  Chemistry 

as  analogous  to  the  white  ash  formed  in  the  burning  of 
magnesium  or  the  black  ash  formed  by  the  burning  of  iron 
filings. 

14.  Heating  the  Red  Ash  of  Mercury. — If  we  take  some  of  the 
red  ash  of  mercury,  place  it  in  a  glass  test  tube,  and  heat  it  very 
strongly  (Fig.  7) ,  we  find  that  it  changes  in  a  remarkable  way : 
first  it  turns  black,  and  then  at  red  heat  it  gradually  grows 
smaller,  until  after  a  few  minutes  none  of  it  remains.  At  the 
same  time,  however,  on  the  cooler  part  of  the 
wall  of  the  tube  a  silvery-looking  coating  has 
appeared,  which  when  the  tube  has  cooled 
may  be  brushed  to  the  bottom  of  the  tube, 
and  is  then  readily  seen  to  consist  of  drops  of 
liquid  mercury.  Thus  by  heating  the  red 
powder  to  a  higher  temperature  than  that  used 
in  its  formation,  mercury  is  reproduced.  But 
this  is  only  half  the  story. 

The  more  important  part  remains  to  be 
told.  Lavoisier  reasoned  about  the  matter 
somewhat  as  follows:  If  burning  substances 
require  air;  if  a  part  of  the  air  disappears  (in 
some  cases  at  least)  during  burning;  if  in  the 
burning  of  metals  like  magnesium  and  iron  the  ash  is  heavier  than 
the  metal  burned ;  if,  as  is  indeed  a  fact,  air  has  weight;  is  it  not 
possible  that  the  burning  substance  unites  with  a  part  of  the  air 
to  form  a  new  kind  of  substance,  and  that  this  new  substance, 
for  example,  magnesium  ash,  is  heavier  than  the  substance 
burned  because  it  contains  not  only  the  latter  but  also  a  part  of 
the  air?  Perhaps  also  the  red  ash  formed  by  the  gentle  heating 
of  mercury  in  contact  with  air  is  also  made  up  of  mercury  and 
something  taken  from  the  air.  Perhaps  the  one-fifth  of  the  air 
that  vanished  has  combined  with  the  mercury  to  form  the  red 
ash.  If  all  these  suppositions  are  true,  perhaps  when  the  red 
ash  was  changed  again  into  mercury  by  being  strongly  heated 
there  was  set  free  at  the  same  time  the  part  of  the  air  which  by 
originally  uniting  with  the  mercury  produced  the  red  ash.  If 
all  this  were  true,  how  could  it  be  proved?  Let  us  see. 


The  Burning  of  Substances — Oxygen  n 

15.  The  Active  Part  of  the  Air:  Oxygen. — The  part  of  the  air 
which  disappeared  may  be  just  that  part  which  causes  substances 
to  burn.     If  it  were  to  be  obtained  pure,  free  from  the  inert  four- 
fifths  which  does  not  support  burning  (combustion),  it  ought  to 
support  combustion  far  better  than  common  air.    This  is  a  matter 
easily  put  to  the  test  of  experiment.    Let  us  again  heat  some 
of  the  red  powder  in  a  test  tube  and  at  the  same  time  thrust  into 
the  tube  a  burning  wood  splint.     We  see  that  it  burns  much  more 
fiercely  and  brightly  than  in  common  air.     Furthermore,  if  we 
have  no  flame,  but  only  a  tiny  spark  on  the  end  of  the  splint, 
we  see  that  when  thrust  into  the  tube  above  the  heated  red  ash 
the  spark  bursts  into  a  vigorous  flame.    The  suppositions  seem 
to  be  true.    Lavoisier  was  led  in  this  way  to  the  discovery  of 
the  secret  of  the  nature  of  burning.    He  called  the  gas  formed 
by  the  heating  of  the  red  powder  oxygen.     This  gas  forms  one 
fifth  by  volume  of  the  air  and  is  the  part  of  the  air  which  is  necessary 
for  the  burning  of  substances.    The  other  four-fifths  by  volume 
of  the  air  is  inert;  it  does  not  support  combustion;  neither  does 
it  support  the  respiration  of  animals.     Lavoisier  called  it  azote; 
we  call  it  nitrogen. 

16.  The  Properties  of  Oxygen. — Oxygen  is  an  invisible  gas 
like  air;  it  has  no  odor  and  it  supports  combustion  far  better  than 

does  air.  By  the  same  method  as  that  employed 
in  the  case  of  air  (chap,  i),  we  may  readily  find 
that  i  liter  of  oxygen  at  a  temperature  of  o°  and 
76  cm.  pressure  weighs  i .  43  g.  It  has,  there- 
fore, a  somewhat  greater  density  than  air,  of 
which  i  liter  weighs  i .  29  g.  Further  evi- 
dence that  the  explanation  of  the  nature  of 
burning,  given  in  the  preceding  paragraph,  is  the 
correct  one  is  furnished  by  experiments  which 
we  may  now  consider. 

17.  Burning  Iron  in  Oxygen. — If  we  place  a 
gram  or  two  of  iron  filings  and  a  minute  piece  of 
phosphorus  on  a  piece  of  asbestos  paper  in  the 
bottom  of  a  3oo-c.c.  round-bottomed  flask  filled  with  pure 
oxygen  and  fitted  with  a  rubber  stopper  and  a  glass  stopcock 


12  Introduction  to  General  Chemistry 

(Fig.  8),  we  shall  find  that  the  weight  of  the  flask  with  its  con- 
tents does  not  change  if  by  heating  we  cause  the  iron  to  burn 
in  the  oxygen.  Now,  we  know  that  when  iron  burns,  the  prod- 
uct weighs  more  than  the  original  iron.  We  know  also  that 
oxygen  has  weight  and  that  the  total  weight  of  the  flask  with  its 
contents  has  not  changed  during  the  burning.  What  then  is  the 
cause  of  the  increase  of  the  weight  of  iron  when  burned?  If  we 
open  the  stopcock  while  its  open  end  is  held  under  water,  we  find 
that  the  water  nearly  fills  the  flask.  We  must  conclude  that  the 
oxygen  has  disappeared.  Is  it  not  reasonable  to  suppose  that 
the  ash  resulting  from  the  burning  of  iron  is  composed  of  the  iron 
originally  taken  and  the  oxygen  which  has  disappeared?  Our 
experiment  has  shown  that  the  weight  of  this  ash  is  precisely  the 
same  as  the  combined  weights  of  the  iron  and  the  oxygen  which 
disappeared  in  the  burning. 

It  will  readily  be  seen  that  the  experiment  with  iron  is  similar 
to  that  made  by  Lavoisier  with  mercury — with  the  difference 
that  iron  burns  rapidly,  whereas  mercury  changes  but  slowly  in 
oxygen.  Furthermore,  the  fact  that  the  red  ash  of  mercury 
when  strongly  heated  gives  again  mercury  and  oxygen  makes 
it  practically  certain  that  the  red  ash  was  formed  by  the  combina- 
tion or  union  of  mercury  with  oxygen  which  composed  part  of 
the  original  air  used  in  Lavoisier's  experiment.  Instead  of  iron, 
in  the  experiment  described,  we  might  have  substituted  magne- 
sium or  phosphorus,  or  indeed  any  one  of  a  large  number  of  other 
substances.  In  each  case  the  result  would  have  been  similar 
to  that  in  the  case  of  iron  and  oxygen  and  a  similar  conclusion 
would  have  been  forced  upon  us.  In  all  such  cases  we  would 
conclude  that  the  process  of  burning  consists  in  the  combination 
or  union  of  gaseous  oxygen  with  the  solid  substance  burned  to  form 
the  product  of  the  combustion. 

18.  Burning  Charcoal  in  Oxygen. — If  we  put  a  piece  of  burn- 
ing charcoal  into  a  bottle  containing  oxygen  we  notice  that  it 
burns  even  more  rapidly  in  oxygen  than  in  the  air.  In  this  case 
there  is  but  a  trifling  amount  of  ash  left  compared  with  the 
amount  of  charcoal  burned.  In  order  to  see  whether  an  invisible 
product  may  have  been  produced  we  may  make  the  following 
experiment.  If  we  pour  a  little  limewater  into  a  bottle  contain- 


The  Burning  of  Substances — Oxygen  13 

ing  oxygen  and  shake  the  limewater  with  the  oxygen  we  notice 
no  change.  If  now  we  pour  limewater  into  a  bottle  in  which 
charcoal  has  been  burned  in  oxygen  and  again  shake  the  con- 
tainer, the  limewater  becomes  milky  in  appearance.  We  must  con- 
clude that  some  invisible  substance,  different  from  oxygen,  has 
been  produced  in  the  latter  case.  If  the  burning  of  charcoal  is 
thought  to  be  analogous  in  nature  to  the  burning  of  iron,  then 
we  might  expect  that  the  product  would  be  something  composed 
of  carbon  and  oxygen  and  that  its  weight  should  be  equal  to  the 
combined  weights  of  the  carbon  burned  and  the  oxygen  taken  up. 
We  can  get  some  evidence  that  this  is  the  case  by  means  of  the 
following  experiment. 

19.  Carbon  Dioxide. — A  small  quantity  of  charcoal  is  placed 
near  one  end  of  a  hard  glass  tube,  the  other  end  of  which  contains 
pieces  of  caustic  soda  (Fig.  9).  If  we  now  weigh  the  tube,  which 


FIG.  9 

may  be  fitted  at  the  end  containing  the  charcoal  with  a  stopper 
and  a  small  glass  tube,  and  then  cause  the  charcoal  to  burn  in  a 
stream  of  oxygen  gas  which  we  may  pass  through  the  tube,  we 
shall  find  that  there  is  an  increase  of  weight,  due  to  the  fact  that 
the  product  formed  by  burning  the  charcoal  has  been  absorbed 
by  the  caustic  soda  in  the  tube.  If  we  place  some  caustic  soda 
in  a  beaker,  dissolve  it  in  water,  and  add  some  hydrochloric  acid, 
we  can  see  no  marked  change.  If  we  treat  the  material  from  the 
charcoal  experiment  in  the  same  way,  we  notice  that  a  gas  is 
given  off  when  we  pour  the  acid  into  the  solution.  A  test  of  this 
gas  with  limewater  shows  that  it  behaves  like  that  obtained 


14  Introduction  to  General  Chemistry 

when  charcoal  is  burned  directly  in  oxygen.  The  results  of 
these  experiments  lead  us  to  conclude  that  when  charcoal  is 
burned  an  nvisible  gas  is  produced,  and  that  this  gas  is  heavier 
than  the  charcoal  burned;  and,  in  fact,  if  charcoal  had  been 
burned  in  a  closed  vessel  with  oxygen,  we  should  find  that  the 
weight  of  vessel  and  contents  had  not  changed  during  the 
burning,  and  would  be  forced  to  conclude  that  the  weight  of 
the  invisible  product  was  just  equal  to  the  sum  of  the  weights  of 
the  charcoal  burned  and  the  oxygen  which  had  united  with  it. 
This  gaseous  product  of  the  burning  of  charcoal  was  formerly 
called  carbonic  acid  gas,  but  is  now  usually  called  carbon  dioxide. 
20.  Experiments  with  a  Burning  Candle. — We  find  by  experi- 
ment that  carbon  dioxide  is  formed  when  wood,  coal,  illuminating 
gas,  gasoline,  etc.,  burn.  We  may  easily  show  by  the  limewater 
test  that  it  is  alsoxformed  during  the  burning  of  a  candle.  We 

may  also  show  that  another  well- 
known  substance  is  produced  when 
the  candle  burns.  If  we  burn  the 
candle  under  an  inverted  funnel 
connected  by  means  of  a  glass 
tube  with  a  U-tube  which  is  cooled 
by  immersion  in  a  vessel  of  mercury 
and  draw  air  through  the  funnel 
and  U-tube  we  find  that  a  colorless 
liquid  collects  in  the  cold  \J-tube 
(Fig.  10).  This  liquid  is  water. 
The  burning  of  the  candle  gives, 

both  carbon  dioxide  and  water.  We  may  readily  show  that 
the  weight  of  the  products  of  a  burning  candle,  if  these  are  suit- 
ably collected,  is  greater  than  the  weight  of  the  candle  burned. 
To  do  this  we  make  use  of  the  arrangement  shown  in  Fig.  n. 
The  candle  is  inclosed  in  a  glass  cylinder,  closed  below  by  a  cork 
having  three  or  four  holes  for  the  admission  of  air.  The  top  of 
the  cylinder  is  filled  with  pieces  of  a  solid  substance  (caustic 
potash)  which  readily  absorbs  both  carbon  dioxide  and  water,  but 
not  oxygen  or  nitrogen,  the  components  of  air.  The  apparatus 
thus  arranged  is  suspended  on  one  side  of  a  balance  and  counter- 
poised. 


The  Burning  of  Substances — Oxygen  15 

The  candle  is  now  lighted  and  allowed  to  burn  for  ten  or 
fifteen  minutes,  whereupon  it  will  be  found  that  the  apparatus  has 
become  appreciably  heavier.  The  increase  in  weight  is  due  to  the 
fact  that  the  carbon  dioxide  and  water  formed  weigh  more  than  the 
candle  burned.  In  fact,  the  excess  weight  is  exactly  the  weight 
of  the  oxygen  which  has  been  consumed  in  the  burning.  Under 
ordinary  circumstances  the  carbon  dioxide  and  water  escape 
our  notice  because  both,  the  latter  being  in  the  form  of  steam, 
are  invisible  gases. 


FIG. 


21.  The  Law  of  the  Conservation  of  Matter.  —  By  the  study 
of  such  facts  as  those  discussed  in  the  preceding  paragraph  and 
many  others  of  a  similar  nature,  Lavoisier  arrived  at  the  con- 
clusion that  when  a  substance  burns  it  unites  with  the  oxygen  of  the 
airt  and  that  the  weight  of  the  product  is  always  exactly  equal  to  the 
weight  of  the  substance  burned  plus  the  weight  of  the  oxygen  which 
unites  with  the  burned  substance  during  the  combustion.  The 
product  may  be  a  solid,  a  liquid,  or  a  gas.  If  it  is  a  volatile 
liquid  or  a  gas  it  usually  escapes  notice  because  it  is  invisible. 
Burning,  therefore,  consists  in  a  union  of  the  substance  burned  with 
oxygen.  In  this  sense  a  substance  which  is  burned  is  not  destroyed; 
the  material  or  matter  composing  it  merely  passes  into  another 
form,  the  quantity  of  matter  in  all  cases  being  measured  by  its 
weight.  These  facts  are  briefly  summed  up  in  the  statement  that 
matter  is  indestructible,  a  statement  which  is  frequently  referred 
to  as  the  Law  of  the  Conservation  of  Matter. 


CHAPTER  III 
PURE   SUBSTANCES— ELEMENTS 

22.  Bodies  and  Substances. — We  use  the  words  "  substance" 
and  "body"  in  chemistry  in  very  definite  senses.  We  speak  of 
things  like  watches  or  knives  as ' '  bodies. ' '  We  say  that  the  blade 
of  the  knife  is  steel,  the  handle  is  pearl.  We  say  that  a  watch  has 
a  case  of  gold  and  a  watch  crystal  of  glass.  We  call  steel,  pearl, 
gold,  and  glass  "substances."  A  substance  is  thus  a  particular 
kind  of  material,  a  body  is  an  object  which  may  be  composed  of 
one  or  many  kinds  of  substances.  Water,  salt,  and  sugar  are 
further  examples  of  substances  in  the  sense  of  this  definition. 


FIG.  12 

23.  Pure  Substances. — We  find  that  natural  waters,  as  those 
of  lakes,  rivers,  and  springs,  are  not  all  alike.  It  now  becomes 
important  to  discover  the  cause  of  the  differences  between 
waters  from  different  sources.  If  we  boil  a  quantity  of  lake 
water  we  find  when  the  water  has  entirely  disappeared  that  a 
solid  residue  is  left.  If  the  steam  from  the  boiling  water  is  con- 
densed by  cooling  it,  as  by  means  of  a  condenser  (Fig.  12)  through 
the  outer  tube  of  which  a  stream  of  cold  water  flows,  we  obtain 
what  is  called  distilled  water.  If  we  now  evaporate  to  dryness 
a  quantity  of  this  distilled  water  we  find  that  no  residue  is  left. 
If  we  prepare  distilled  water  from  any  natural  water  we  find  that 

16 


Pure  Substances — Elements 


it  will  always  evaporate  completely,  leaving  no  solid  residue. 
We  find  further  that  different  kinds  of  natural  water  leave  differ- 
ent proportions  of  solid  residue  upon  evaporation  and  that  the 
nature  of  the  solid  material  left  also  differs  in  different  cases,  but 
that  the  distilled  water  in  one  case  cannot  be  distinguished  in 
any  way  from  that  obtained  in  another.  We  say  then  that 
distilled  water  is  pure  water,  a  pure  substance,  and  the  natural 
waters  are  not  pure  water,  but  that  they  contain  dissolved  foreign 
substances.  If  the  natural  water  is  muddy,  that  is,  if  it  is  not 

clear,  the  foreign  material 
which  causes  it  to  appear 
muddy  can  be  separated 
by  filtration  (Fig.  13),  a 
process  in  which  the  liquid 
is  allowed  to  seep  through 
a  piece  of  filter  paper 
folded  so  as  to  fit  snugly 
into  a  funnel.  The  mud 
remains  on  the  filter  paper. 
However,  filtration  will  not 


FIG.  13 


remove  any  of  the  dissolved  material,  but  only  that  which  is 
suspended  in  the  water. 

24.  Pure  Salt  Made  from  Rock  Salt. — Common  salt  is  found 
in  nature  as  a  mineral  known  as  rock  salt.  We  find  that  different 
samples  of  rock  salt  differ  in  color,  taste,  specific  gravity,  and  in 
other  ways.  If  we  mix  rock  salt  with  water  we  find  that  a  large 
part  dissolves  in  the  water.  In  general  a  small  amount  of 
material,  sand,  etc.,  will  not  dissolve,  even  though  we  take  a  large 
amount  of  water.  If  we  filter  the  solution  we  separate  the  water 
and  dissolved  material  from  the  part  which  has  not  dissolved. 
That  which  runs  through  the  filter  paper  is  called  the  filtrate ;  it 
is  the  solution  of  the  salt  in  water.  If  we  boil  away  the  water 
we  find  that  the  salt  is  left  in  the  solid  form  and  that  the  material 
is  now  free  from  color,  that  is,  that  it  is  white,  and  that  it  will 
dissolve  completely  in  water.  The  salt  so  prepared  is  purer  than 
the  rock  salt  taken.  Just  as  it  is  possible  to  prepare  pure  water 
from  any  natural  water,  so  analogously  it  is  possible  to  prepare 


i8 


Introduction  to  General  Chemistrv 


pure  salt  from  any  natural  salt.  Pure  salt  is  always  exactly  the 
same  in  taste,  color,  specific  gravity,  etc.,  from  whatever  source 
it  may  have  been  obtained.  The  process  for  the  purification 
of  salt,  described  in  the  statement  above,  gives  in  all  cases  a 
much  purer  product  than  the  original  rock  salt — pure  enough 
for  table  use,  but  not  a  perfectly  pure  substance.  It  still  con- 
tains very  small  amounts  of  some  foreign  substances;  but  even 
these  can  be  removed  by  well-known  methods  which  the  student 
will  learn  later.  A  pure  substance  is  a  substance  which  consists 
of  one  sort  of  material.  It  always  has  definite  physical  proper- 
ties, from  whatever  source  it  may  be  obtained. 

25.  Decomposition  of  Substances. — It  was  found  that  the  red 
ash  formed  by  heating  mercury  in  contact  with  air  was  changed, 
upon  being  heated  still  more,  into  mercury  and  oxygen.  We  say 
in  this  case  that  the  red  ash  of  mercury  has  been  decomposed  into 

mercury  and  oxygen.  We  can 
accomplish  the  decomposition  of 
many  substances  in  an  equally 
simple  fashion. 

We  will  now  consider  a  few 
such  cases  as  illustrations. 

26.  Decomposition  of  Sal 
Soda. — If  we  place  in  a  test  tube 
a  crystal  of  common  washing-soda, 
also  known  as  sal  soda,  and  heat  it 
gently  over  a  Bunsen  flame,  we  find 
that  water  is  produced  as  steam, 
and  that  it  condenses  in  the  cold 
end  of  the  test  tube.  An  opaque 
solid  is  left  in  place  of  the  clear 
crystal  of  sal  soda  taken.  We  say 
that  the  sal  soda  has  been  decom- 
posed into  dry  soda  and  water.  It 
would  be  easy  to  show  that  the  weight  of  the  water  and  dry  soda 
formed  is  equal  to  the  weight  of  sal  soda  taken.  In  other 
words,  the  sal  soda  has  been  decomposed  into  dry  soda  and 
water. 


FIG.  14 


Pure  Substances — Elements  19 

27.  Electrolysis  of  Water. — If  we  pass  an  electric  current 
through  some  water  (Fig.  14)  to  which  we  nave  added  a  few  drops 
of  sulfuric  acid,  we  find  that  gases  are  produced  at  the  platinum 
electrodes.     The  decomposition  of  a  substance  by  an  electric 
current  is  called  electrolysis.    If  we  collect  each  of  these  gases 
separately  we  find  that  one  of  them  is  oxygen.     The  other  gas, 
the  volume  of  which  is  double  that  of  the  oxygen,  has  quite  differ- 
ent properties;    it  is  called  hydrogen.     If  we  bring  a  lighted 
splinter  into  the  oxygen,  the  splinter  continues  to  burn  with 
increased  brilliancy  and  rapidity.     If  we  repeat  this  test  with 
hydrogen,  we  find  that  the  hydrogen  itself  catches  fire,  just  as 
illuminating  gas  would  do,  and  that  the  splinter  itself  no  longer 
burns  in  the  hydrogen  gas.     These  facts  may  be  concisely  stated 
by  saying  that  oxygen  supports  combustion,  while  hydrogen  burns 
but  does  not  support  combustion.    It  would  be  possible  to  show 
by  experiment  that  the  weight  of  the  water  decreases  during  the 
passage  of  the  electric  current  through  it,  and  that  this  decrease 
in  weight  is  just  equal  to  the  combined  weights  of  the  oxygen  and 
hydrogen  formed.    The  total  amount  of  sulfuric  acid  added  to 
the  water  remains  in  the  water  at  the  end  of  the  electrolysis  and 
would  serve  to  promote  the  decomposition  of  any  desired  amount 
of  water.    The  complete  explanation  of  the  behavior  of  the 
sulfuric  acid  cannot  be  given  at  this  point,  but  we  know  that  the 
hydrogen  and  oxygen  formed  come  exclusively  from  the  water 
and  not  from  the  acid  nor  the  platinum  nor  the  glass  of  the 
vessel  used.    We  conclude  that  water  is  decomposed  by  the  electric 
current  into  hydrogen  and  oxygen.    Therefore  we  may  say  that 
water  is  composed  of  hydrogen  and  oxygen  or  that  water  is  a 
compound  of  hydrogen  and  oxygen.    As  a  matter  of  fact,  when 
hydrogen  burns  in  air  water  is  formed.     If  a  cold  beaker  is  held 
over  a  jet  of  burning  hydrogen,  water  will  be  seen  to  condense  in 
a  mist  on  the  surface  of  the  beaker. 

28.  Magnesium  Burned  in  Steam. — That  water  is  composed 
of  oxygen  and  hydrogen  may  be  shown  in  many  other  ways,  one 
of  which  is  the  following.    When  a  piece  of  magnesium  ribbon 
burns  in  air  the  magnesium  unites  with  the  oxygen  of  the  air 
to  form  a  white  solid  which  we  call  magnesium  oxide.    Now, 


20 


Introduction  to  General  Chemistry 


magnesium  will  also  burn  in  steam  (Fig.  15)  nearly  as  readily 
as  it  does  in  the  air  or  in  pure  oxygen,  and  we  find  that  the 
white  solid  which  is  again  formed  is  also  magnesium  oxide. 
In  addition,  hydrogen  gas  is  produced  and  may  easily  be  col- 
lected over  water.  Since  magnesium  oxide  is  composed  of 
magnesium  and  oxygen,  and  we  obtain  from  magnesium  and 
water  magnesium  oxide  and  hydrogen,  we  are  again  led  to  the 
conclusion  that  water  is  composed  of  hydrogen  and  oxygen. 

29.  Steam  Passed  over  Hot  Iron. — An  entirely  analogous 
experiment  may  be  carried  out  with  iron  and  steam.     In  this 


FIG.  15 

case  iron  turnings  or  fine  iron  wire  is  strongly  heated  in  an  iron  or 
glass  tube  (Fig.  16).  When  steam  is  passed  through  the  tube, 
iron  oxide  and  hydrogen  are  produced,  a  result  which  leads  to  the 
same  conclusion  as  before  regarding  the  composition  of  water. 

30.  Magnesium  Burned  in  Carbon  Dioxide. — The  composi- 
tion of  carbon  dioxide  may  be  discovered  by  burning  magnesium 
in  this  gas.  We  find  that  magnesium  oxide  and  a  product 
resembling  charcoal  are  formed.  The  latter  substance  is  carbon, 
of  which  charcoal  is  a  nearly  pure  form.  We  conclude,  therefore, 
that  carbon  dioxide  is  composed  of  carbon  and  oxygen  or  is  a 
compound  of  carbon  and  oxygen. 

The  facts  already  considered  lead  to  the  conclusion  that  the 
red  ash  obtained  when  mercury  is  heated  gently  in  air  is  com- 


Pure  Substances — Elements 


21 


posed  of  mercury  and  oxygen;  briefly,  that  it  is  a  compound  of 
mercury  and  oxygen — a  fact  represented  by  the  chemical  name 
of  the  red  ash,  mercuric  oxide. 

31.  Elements. — The  substances  mercury,  oxygen,  hydrogen, 
and  carbon  have  never  been  decomposed  into  simpler  substances. 
We  say  that  hydrogen  and  oxygen  are  the  elements  of  which 
water  is  composed;  that  carbon  and  oxygen  are  the  elements  com- 
posing carbon  dioxide. 

We  may  discover  of  what  elements  a  substance  is  composed  in 
two  ways:  either  by  the  decomposition  of  the  substance  into  the 


FIG.  1 6 

simpler  ones  which  compose  it — the  process  called  analysis,  or 
by  causing  known  elements  to  unite  in  the  formation  of  the 
original — the  process  called  synthesis.  As  a  result  of  the 
electrolysis  of  water  we  have  concluded  that  water  is  composed 
of  hydrogen  and  oxygen.  This  conclusion  may  now  be  tested 
by  seeing  whether  water  can  be  obtained  from  hydrogen  and 
oxygen.  We  found  that  hydrogen  burns  readily.  If  we  burn 
a  jet  of  hydrogen  under  an  inverted  funnel  and  draw  the  product 
through  a  cooled  U-tube,  as  in  the  experiment  with  the  candle, 
we  shall  find  that  liquid  water  collects  in  the  U-tube  and  that 
the  most  careful  search  fails  to  reveal  any  other  substance  as 
the  product  of  the  burning  of  hydrogen  in  air  or  in  pure  oxygen. 
Water  is,  therefore,  a  compound  of  the  elements  hydrogen  and 


22  Introduction  to  General  Chemistry 

oxygen.  Since  the  burning  of  charcoal,  which  is  a  nearly  pure 
form  of  the  element  carbon,  gives  carbon  dioxide  and  nothing 
else,  we  know  that  carbon  dioxide  is  a  compound  of  the  elements 
carbon  and  oxygen. 

32.  The  Burning  of  Copper;  Copper  Oxide. — If  the  metal 
copper,  in  the  form  of  fine  wire  or  filings,  is  heated  in  air  or  in 
oxygen,  it  is  slowly  changed  into  a  black  substance  quite  different 
in  appearance  from  metallic  copper;  but  during  this  change  we 
do  not  observe  the  production  of  any  light.     By  means  of  the 
balance  we  may  find  that  the  black  substance  formed  is  heavier 
than  the  copper  taken,  and  we  at  once  suspect  that  the  copper 
has  united  with  oxygen  to  form  a  compound.     If  the  heating 
of  the  copper  had  been  carried  out  in  a  sealed  glass  vessel  con- 
taining oxygen,  as  in  the  earlier  experiment  with  iron  powder, 
it  would  have  been  found  that  gaseous  oxygen  had  disappeared 
and  that  the  weight  of  the  black  product  was  exactly  equal  to  the 
weight  of  the  copper  taken  plus  the  weight  of  the  gaseous  oxygen 
which  had  disappeared.    The  black  substance  would  seem,  there- 
fore, to  be  a  compound  of  copper  and  oxygen.    We  know  that 
when  the  red  mercury  oxide  is  strongly  heated  it  is  decomposed 
into  mercury  and  oxygen.     If  we  heat  the  black  product  from 
copper  to  the  highest  temperature  we  can  obtain  with  the  Bunsen 
burner,  we  find  that  it  remains  unaltered  in  weight  and  appear- 
ance and  that  no  oxygen  is  given  off.     This  fact  might  lead  us 
to  suspect  that  the  black  substance  is  not  a  compound  of  copper 
and  oxygen,  since  its  behavior  is  not  analogous  to  that  of  mercury 
oxide.     In  this  connection  the  following  experiment  will  prove 
of  interest. 

33.  Hydrogen  Passed  over  Hot  Copper  Oxide. — If  we  put  two 
or  three  grams  of  the  black  copper  product  in  a  porcelain  boat  in 
a  "hard"  or  difficultly  fusible  glass  tube,  heat  the  tube  and  con- 
tents by  means  of  a  Bunsen  flame,  and  then  pass  a  current  of 
hydrogen  through  the  tube,  we  observe  that  the  solid  glows  or 
seems  to  burn  (Fig.  17).    At  the  same  time  we  notice  that  liquid 
water  condenses  in  the  colder  part  of  the  glass  tube.    After  a  few 
minutes  the  glow  has  disappeared,  even  though  the  stream  of 
hydrogen  has  continued.    At  this  point  the  heating  may  be 


Pure  Substances — Elements  23 

discontinued  and  the  solid  which  is  left  in  the  boat  allowed  to 
cool  in  the  stream  of  hydrogen  gas.  We  now  observe  that  the 
solid  has  the  appearance  and  properties  of  metallic  copper,  which 
in  fact  it  is.  However,  the  copper  is  not  in  a  single  compact 
lump,  for  a  reason  which  must  be  evident.  Metallic  copper  can 
be  melted,  but  the  melting-point  is  a  very  much  higher  tempera- 
ture than  that  attained  in  the  preceding  experiment.  Only  by 
heating  the  copper  to  a  point  above  this  melting  temperature 
could  the  material  be  obtained  in  a  single  lump.  This  could 
easily  be  accomplished  by  directing  an  intense  blowpipe  flame 
upon  the  metal  particles  contained  in  the  porcelain  boat. 

We  may  now  consider  the  nature  of  the  changes  which  occured 
in  this  experiment.     Since  we  obtained  water  and  copper,  and 


FIG.  17 

since  wre  know  that  water  is  a  compound  of  the  element  hydrogen 
with  oxygen,  we  conclude  that  the  oxygen  was  originally  united 
with  the  copper  and  that  the  black  substance  must  have  been  a 
compound  of  copper  and  oxygen.  This  substance  is  called 
copper  oxide.  We  might  express  the  result  in  the  following 
simple  fashion:  Copper  Oxide + Hydrogen ->Water-f  Copper;  or 
instead  of  "  Water"  we  might  write  "Hydrogen  Oxide,"  the  true 
chemical  name  for  water.  This  statement  would  then  show  at 
a  glance  the  nature  of  the  chemical  change  which  had  occurred. 
34.  Discovery  of  the  Elements  Composing  a  Substance :  the 
Analysis  of  Malachite. — There  is  an  almost  innumerable  variety 
of  bodies  on  and  in  the  earth,  but  these  are  composed  of  a  very 
much  smaller  number  of  definite  chemical  substances.  However, 
the  number  of  definite  substances  is  still  very  great,  many 
thousands  having  been  carefully  described.  Chemistry  has  for 
its  object  the  systematic  study  of  pure  substances,  their  properties, 


24  Introduction  to  General  Chemistry 

and  their  behavior  toward  one  another.  Happily  the  study  of  this 
immense  number  of  substances  is  greatly  simplified  by  the  fact 
that  they  are  all  made  up  of  a  relatively  small  number  of  elements. 
The  way  in  which  the  elements  composing  a  substance  of  un- 
known composition  are  discovered  may  be  illustrated  by  means 
of  an  experiment  with  the  mineral  known  as  malachite.  Mala- 
chite is  a  beautiful  crystalline  substance  often  used  as  an  orna- 
mental stone  and  also  as  one  of  the  sources  from  which  a  familiar 
metal  is  obtained.  If  we  place  in  a  test  tube,  fitted  with  a  cork 
and  a  bent  glass  tube,  a  few  grams  of  malachite  and  heat  the 
substance  gently  in  a  flame,  we  notice  that  a  change  in  color 
from  green  to  black  occurs  and  at  the  same  time  that  water 
condenses  in  the  colder  part  of  the  glass  tube  and  a  gas  is  also 
given  off.  If  we  pass  this  gas  into  limewater  we  find  that  it 
behaves  like  carbon  dioxide,  which  in  fact  it  is.  By  means  of 
the  balance  we  might  find  that  the  combined  weights  of  the 
carbon  dioxide,  water,  and  black  product  equal  the  weight  of 
the  original  malachite.  Since  we  know  of  what  elements  carbon 
dioxide  and  water  are  composed,  it  only  remains  to  find  the 
composition  of  the  black  substance  in  order  to  have  a  complete 
knowledge  of  the  elements  composing  malachite.  If  this  black 
substance  were  heated  in  a  stream  of  hydrogen,  it  would  be  found 
to  yield  water  and  a  red  metallic-looking  substance  which  could 
easily  be  recognized  as  copper.  Therefore,  the  black  substance 
must  have  been  copper  oxide.  The  results  may  then  readily 
be  interpreted.  Malachite  when  heated  is  decomposed  into 
carbon  dioxide,  water,  and  copper  oxide.  Knowing  as  we  do  the 
elements  composing  each  of  these  three  products,  we  are  led  to 
the  conclusion  that  malachite  is  a  compound  of  the  elements 
carbon,  oxygen,  hydrogen,  and  copper.  Chemists  have  so  far  been 
unable  to  decompose  copper  into  anything  simpler.  It  is, 
therefore,  known  as  an  elementary  or  simple  substance,  and  we 
say  that  malachite  is  a  compound  of  the  four  elements,  carbon, 
hydrogen,  oxygen,  and  copper. 

35.  Some  Common  Elements. — The  total  number  of  known 
elements  is  about  eighty-five,  of  which  less  than  thirty  are  com- 
mon. In  the  following  partial  list  of  commoner  elements,  the 


Pure  Substances — Elements  25 

student  will  find  the  names  of  ten  or  twelve  familiar  metals. 
Carbon  and  sulfur,  which  are  well  known  to  everyone,  are  not 
metals;  they  are  classed  as  non-metals. 

A  FEW  COMMON  ELEMENTS 

Silver  Copper  Nickel  Carbon 

Gold  Lead  Magnesium  Sulfur 

Platinum  Tin  Zinc  Oxygen 

Iron  Aluminum  Mercury  Hydrogen 


CHAPTER  IV 
THE  LAW  OF  DEFINITE  COMPOSITION 

36.  The    Percentage    Composition    of    Water. — We    have 
already  seen  that  when  an  electric  current  was  passed  through 
water,  the  latter  was  decomposed  into  two  gases,  hydrogen  and 
oxygen.     It  was  found  that  the  volume  of  the  hydrogen  was 
double  that  of  the  oxygen  obtained  in  the  electrolysis.     This  was 
not  a  matter  of  accident,  for  it  is  always  found  that  the  same 
result  is  obtained  whenever  water  is  electrolyzed.     Since  water 
is  composed  only  of  hydrogen  and  oxygen,  we  may  calculate  the 
percentages  of  hydrogen  and  oxygen  by  weight  if  we  know  the 
weight  of  a  liter  of  each  of  these  gases.    Direct  weighing  of  the 
gases  has  shown  that  i  liter  of  hydrogen  weighs  o .  090  g.  and 
i  liter  of  oxygen  weighs  1.429  g.,  the  gases  being  weighed  at  o° 
and  76  cm.  pressure.     From  these  figures  it  is  easy  to  calculate 
that  water  is  composed  of  11.2  per  cent  of  hydrogen  and  88 . 8 
per  cent  of  oxygen  by  weight.     Pure  water  prepared  from  any 
source  whatever  always  has  exactly  this  composition. 

The  percentage  composition  of  water  may  also  be  found  in 
another  way.  It  was  found  in  section  33  that  water  and  copper 
are  formed  when  hydrogen  is  passed  over  heated  copper  oxide. 
If  this  experiment  be  carried  out  with  a  weighed  quantity  of 
copper  oxide,  and  the  weight  of  copper  which  remains  after  the 
experiment  is  found,  the  difference  in  the  two  weights  will  repre- 
sent the  weight  of  oxygen  contained  in  the  water  which  has  been 
formed.  If  the  weight  of  the  water  is  determined,  then  the 
percentage  of  oxygen  in  water  may  readily  be  calculated.  In 
this  case  we  find  precisely  the  same  result  as  that  given  in  the 
preceding  paragraph. 

The  details  of  the  experiment  are  as  follows. 

37.  The  Quantitative  Synthesis  of  Water. — About  one  gram 
of  pure  copper  oxide  is  placed  in  a  weighed  porcelain  boat  and 
heated  sufficiently  to  drive  off  the  moisture  which  it  may  con- 

26 


The  Law  of  Definite  Composition  27 

tain.1  The  boat  and  contents  are  weighed  as  soon  as  cool  and 
placed  at  once  in  a  hard  glass  tube.  This  tube  (Fig.  18)  is  con- 
nected at  each  end  with  U- tubes  filled  with  calcium  chloride,  a 
substance  that  absorbs  water  with  great  readiness.  One  of 
these  U- tubes  is  connected  with  a  source  of  hydrogen  gas  and 
serves  to  remove  all  moisture  (water  vapor)  from  the  hydrogen. 
The  other  U-tube  will  serve  to  absorb  the  water  formed  in  the 
chemical  reaction  between  the  copper  oxide  and  the  hydrogen. 


FIG.  18 

This  second  U-tube  is  accurately  weighed  at  the  beginning  of 
the  experiment. 

When  all  is  ready,  the  stream  of  hydrogen  is  started  and  con- 
tinued until  all  air  is  driven  from  the  tubes.  The  tube  contain- 
ing the  boat  is  now  heated  until  the  reaction  begins,  and  kept  hot 
enough  beyond  the  boat  to  prevent  the  condensation  of  the 
steam  formed,  which  is  carried  by  the  stream  of  hydrogen  into 
the  weighed  U-tube. 

When  all  the  copper  oxide  has  been  changed  into  copper  and 
the  water  has  all  been  driven  over  into  the  U-tube,  the  heating 
is  discontinued  and  the  copper  allowed  to  cool  in  a  stream  of 
hydrogen.  The  hydrogen  is  then  driven  out  by  a  stream  of  air, 
and  the  U-tube  detached  and  weighed.  The  object  in  repla- 
cing the  hydrogen  by  air  is  readily  understood  when  one  recalls 
that  hydrogen  is  far  lighter  than  air.  Therefore  the  weight  of 
the  tube  filled  with  hydrogen  would  be  appreciably  less  than 
if  it  is  filled  with  air.  The  increase  in  weight  is  the  weight 
of  the  water  formed.  The  boat  containing  the  copper  is  also 
weighed.  The  loss  in  weight  is  the  weight  of  oxygen  contained 

1  Most  substances,  especially  if  porous  or  in  the  form  of  powder,  absorb  more 
or  less  moisture  from  the  air. 


28  Introduction  to  General  Chemistry 

in  the  water  formed.    The  results  of  an  actual  experiment 
were  as  follows: 

Boat  and  copper  oxide 9 . 523  g. 

Boat 8.451 

Copper  oxide i .  072  g. 

Boat  and  copper 9.311  g. 

Boat 8.451 

Copper o .  860  g. 

Tube  and  water 18 . 665  g. 

Tube 18.426 


Water o ,  239  g. 

Since  1.072  g.  —  0.860  g.  =  0.21 2  g.,  we  conclude  that  o.239g. 
of  water  was  formed  from  o .  2 1 2  g.  of  oxygen,  which  at  the  begin- 
ning was  in  combination  with  the  copper  in  the  form  of  copper 
oxide.  Therefore  water  consists  of  0.212  g.^-o.  239  g.  =  0.887 
=  88.7  per  cent  oxygen.  The  difference  between  the  weight 
of  water  formed  and  that  of  the  oxygen  used  is  the  weight  of 
hydrogen,  which  is  o .  239  g.  —  o .  2 1 2  g.  =  o .  02  7  g.  This  is  readily 
found  to  be  11.3  per  cent  of  the  weight  of  the  water.  Very 
carefully  performed  experiments,  made  in  this  way,  show  that 
water  contains  88 . 8  per  cent  by  weight  of  oxygen  and  11.2  per 
cent  of  hydrogen;  the  difference  of  o.i  per  cent  between  the 
values  found  in  the  lecture  experiment  quoted  and  those  ob- 
tained in  the  most  accurate  experiments  made  by  skilled  chemists 
working  with  greatest  care  and  under  ideal  conditions  is  due  to 
the  experimental  errors  in  the  rather  crude  lecture  experiment. 

38.  The  Percentage  Composition  of  Copper  Oxide. — It  is  also 
easy  to  see  that  we  may  find  the  percentage  composition  of  copper 
oxide  from  the  data  just  considered.  Thus  1.072  g.  of  copper 
oxide  gave  o .  860  g.  of  copper  by  loss  of  o .  2 1 2  g.  of  oxygen;  from 
which  we  find  that  copper  oxide  is  composed  of  80.2  per  cent 
copper  and  19.8  per  cent  oxygen.  The  most  accurate  experi- 
ments made  in  this  way  give  79.9  per  cent  copper  and  20.1 
per  cent  oxygen,  the  difference  being  due  to  experimental  error 
in  the  lecture  experiment.  Pure  copper  oxide  always  has  exactly 
the  composition  shown  by  these  figures. 


The  Law  of  Definite  Composition  29 

39.  The  Percentage  Composition  of  Carbon  Dioxide. — We 
have  found  that  carbon  in  the  form  of  charcoal  burns  readily  in 
air  or  in  oxygen  with  the  formation  of  a  colorless  gas  called  carbon 
dioxide.  The  percentage  composition  of  carbon  dioxide  may 
be  found  by  burning  a  known  weight  of  pure  carbon  in  oxygen 
gas  and  rinding  the  weight  of  carbon  dioxide  formed.  It  will  be 
recalled  that  carbon  dioxide  is  easily  absorbed  by  solid  caustic 
soda.  It  is  also  readily  absorbed  by  a  solution  of  caustic  potash 
in  water,  while  neither  oxygen  nor  air  is  absorbed  by  such  a  solu- 
tion. If  the  gases  formed  by  the  burning  of  carbon  in  a  stream 
of  oxygen  are  passed  through  a  suitable  bulb  containing  caustic 
potash  solution,  all  of  the  carbon  dioxide  will  be  retained  by  the 
solution  and  the  oxygen  will  pass  through  unabsorbed.  The 
increase  in  weight  of  the  bulb  will  represent  the  weight  of  the 
carbon  dioxide  formed  by  the  burning  of  the  carbon. 


FIG.  19 

The  arrangement  of  the  apparatus  is  shown  in  Fig.  19. 
About  o .  2  g.  of  pure  carbon,  made  from  sugar,  is  contained  in  a 
porcelain  boat  which  is  placed  in  a  hard  glass  tube  connected  at 
one  end  with  a  supply  of  pure  oxygen  and  at  the  other  with  a 
calcium  chloride  tube  and  a  weighed  potash  bulb,  which  contains 
a  30  per  cent  solution  of  caustic  potash.  The  middle  part  of  the 
tube  should  contain  a  column  of  copper  oxide,  to  insure  the 
complete  conversion  of  the  carbon  into  carbon  dioxide.  The 
calcium  chloride  tube  serves  to  catch  any  moisture  present. 
The  carbon  is  ignited  by  heating  the  tube  with  a  gas  burner; 
after  the  carbon  has  completely  burned  and  all  of  the  carbon 
dioxide  formed  has  been  driven  over  into  the  potash  bulb  by 
the  stream  of  oxygen,  a  slow  stream  of  air  is  blown  or  drawn 
through  the  apparatus  to  replace  the  oxygen  by  air.  The 
potash  bulb  is  then  detached  and  weighed.  In  an  actual  lecture 
experiment  o.  194  g.  of  carbon  yielded  o.  701  g.  of  carbon  dioxide; 


30  Introduction  to  General  Chemistry 

from  which  we  find  that  this  gas  contains  27.6  per  cent  of  carbon 
and  72.4  per  cent  of  oxygen0  The  most  accurate  experiments 
of  skilled  chemists  show  the  correct  percentages  to  be  27.3  per 
cent  carbon  and  72.7  per  cent  oxygen. 

40.  The  Action  of  Sodium  on  Water:  Caustic  Soda. — It  is  a 
matter  of  importance  to  know  the  exact  percentage  composition 
of  pure  substances  and  a  great  variety  of  methods  must  be 
employed   in   the  making   of   such   determinations.     It  often 
happens  that  the  method  which  would  seem  to  be  most  direct 
and  desirable  is  not  practicable  because  the  violence  of  the  inter- 
action of  the  elements  which  we  bring  together  would  cause  loss 
of  some  of  the  material  taken.     This  may  be  illustrated  by  an 
experiment  with  the  element  sodium.     If  we  throw  a  piece  of 
this  metal  upon  water,  we  observe  that  the  action  is  a  violent 
one  which  ordinarily  ends  in  an  explosion  that  throws  part  of 
the  substance  out  of  the  beaker  in  which  it  was  contained.     We 
may  carry  out  the  same  reaction  without  loss  of  material  and 
obtain  precisely  the  same  product  if  the  piece  of  sodium  is 
exposed  to  water  vapor  instead  of  being  thrown  upon  liquid 
water.     In  this  case  the  reaction  requires  much  more  time,  but 
it  proceeds  quietly  and  without  loss  of  material.     The  white 
solid  so  obtained  is  caustic  soda. 

41.  The  Action  of  Hydrochloric  Acid  on  Caustic  Soda:  Com- 
mon Salt. — If  we  add  to  a  solution  of  caustic  soda  contained  in  a 
beaker  a  sufficient  amount  of  pure  hydrochloric  acid  and  evap- 
orate the  resulting  solution  to  dryness,  we  find  that  the  product 
is  one  with  which  we  are  well  acquainted.     It  is  nothing  more 
nor  less  than  common  salt,  and  if  the  materials  used  are  all  pure 
the  product  will  be  chemically  pure  salt.    We  discover  in  this 
way  that  the  metal  sodium  is  one  of  the  constituents  of  common 
salt.     In  fact,  metallic  sodium  may  be  obtained  by  the  elec- 
trolysis of  molten  salt,  although  this  is  not  the  most  satisfactory 
method  of  making  this  metal.     The  percentage  of  sodium  in  salt 
may  readily  be  found  if  the  weights  of  sodium  taken  and  of  salt 
obtained  are  determined. 

42.  The  Percentage  of  Sodium  in  Common  Salt.— In  an 
actual  experiment  0.483  g.  of  metallic  sodium  was  weighed  in  a 


The  Law  of  Definite  Composition  31 

stoppered  test  tube  (to  prevent  action  of  the  moisture  of  the  air). 
The  sodium  was  placed  on  a  strip  of  silver  foil  which  rested  on  the 
edges  of  a  small  porcelain  dish  containing  about  10  c.c.  of  water, 
and  covered  with  a  beaker.  In  the  course  of  a  few  hours  the 
sodium  had  reacted  completely  with  the  water  vapor  to  form  a 
solution  of  caustic  soda  which  dripped  into  the  dish.  A  little 
of  the  solution  adhering  to  the  foil  was  rinsed  into  the  dish  with 
a  little  water.  Sufficient  pure  hydrochloric  acid  was  then  added 
and  the  solution  evaporated  by  steam  heat  in  the  manner  shown 
in  Fig.  20.  The  beaker  contained  ordinary  water.  By  this 
mode  of  evaporation  of  the  solution  in  the  dish  we  avoid  loss  by 
spattering  that  would  occur  if  we  should 
boil  the  solution  by  heating  the  dish 
directly  with  the  flame.  When  the  salt 
appeared  to  be  dry,  the  dish  was  heated 
very  cautiously  with  the  direct  flame,  to 
drive  off  the  small  amount  of  remaining 
water.  When  cold,  the  dish  and  contents 
were  weighed.  It  was  found  in  this  way  — 


that  0.483  g.  of  sodium  gave  1.217  g-  °f  FIG.  20 

common    salt,  which  indicated   that  salt 
contains  39.7  per  cent  of  sodium.     The  correct  result  is  30.4 
per  cent. 

43.  The  Electrolysis  of  Hydrochloric  Acid:  Chlorine. — It  is, 
of  course,  obvious  that  the  sodium  in  common  salt  must  be  com- 
bined with  one  or  more  elements  and  the  student  will  readily 
guess  that  a  clue  to  the  other  constituents  of  common  salt  may 
be  gained  by  a  knowledge  of  the  constituents  of  hydrochloric 
acid.  If  we  pass  an  electric  current  through  a  concentrated  solu- 
tion of  hydrochloric  acid  contained  in  the  apparatus  shown  in 
Fig.  21,  we  find  that  two  gaseous  products  are  obtained,  the 
volumes  of  which  are  practically  equal.  One  of  these  is  colorless. 
It  is  lighter  than  air  and  burns  with  a  hot  but  non-luminous 
flame  and  in  so  doing  yields  water;  these  properties  show  the 
colorless  gas  to  be  hydrogen.  The  other  gas  is  pale  yellow  in 
color;  it  is  heavier  than  air,  one  liter  weighing  3. 22  g.,  and  has 
an  exceedingly  disagreeable,  irritating  odor.  This  gas  is  known 


Introduction  to  General  Chemistry 


as  chlorine.    Inasmuch  as  chlorine  has  never  been  separated 
into  simpler  substances,  we  conclude  that  it  is  an  element. 

44.  The  Union  of  Hydrogen  and  Chlorine:  Hydrogen 
Chloride  Gas. — Since  the  hydrochloric  acid  which  was  elec- 
trolyzed  contained  water,  we  should  not  be  warranted  in  con- 
cluding that  hydrogen  is  a  constituent  of  hydrochloric  acid; 

for,  as  we  know,  hydrogen  is  also  one  of 
the  constituents  of  water.  If  we  bring 
together  equal  volumes  of  the  gases 
hydrogen  and  chlorine  and  allow  them 
to  mix,  and  if  we  allow  the  vessel  to 
stand  in  diffused  light  for  a  day  or 
two,  we  notice  that  the  yellow  color 
of  the  chlorine  has  disappeared.  We 
find  that  a  colorless  gas  remains  which 
dissolves  with  the  greatest  ease  in 
water,  and  that  neither  hydrogen  nor 
chlorine  is  left.  Since  water  which  has 
y  L  dissolved  the  gas  has  all  of  the  proper- 

ties of  a  solution  of  pure  hydrochloric 
acid,  we  interpret  the  results  as  show- 
ing that  equal  volumes  of  hydrogen 
and  chlorine  gases  combine  to  form 
a  new  gas  which  we  call  hydrogen 
chloride  gas,  and  that  the  latter  when 
dissolved  in  water  constitutes  hydro- 
chloric acid.  Hydrogen  chloride  gas  may  be  distinguished  from 
the  other  gases  which  we  have  met  in  several  ways,  notably  by 
its  marked,  choking  odor,  by  the  fact  that  it  fumes  or  gives  a 
white  cloud  in  moist  air,  and  it  dissolves  with  great  ease  in 
water,  as  well  as  in  several  other  ways. 

45.  Salt  a  Compound  of  Sodium  and  Chlorine. — The  fact  that 
hydrochloric  acid  is  known  to  be  a  compound  of  chlorine  suggests 
that  common  salt  may  also  contain  this  element.  This  is  in 
fact  the  case.  It  can  readily  be  shown  by  experiment  that 
common  salt  results  from  the  union  of  chlorine  gas  with  metallic 
sodium.  Inasmuch  as  nothing  else  is  needed  and  no  other 


FIG.  21. — Brownlee's 
apparatus. 


The  Law  of  Definite  Composition  33 

product  than  salt  is  formed,  we  must  conclude  that  salt  is  a 
compound  of  the  elements  sodium  and  chlorine.  This  fact  is 
indicated  by  the  chemical  name  of  common  salt,  sodium 
chloride.  Since  salt  contains  39.4  per  cent  of  sodium,  the 
percentage  of  chlorine  must  be  60.6. 

46.  The  Law  of  Definite  Composition. — The  preceding  para- 
graphs of  this  chapter  are  intended  to  illustrate  how  we  may 
arrive  at  a  knowledge  of  the  nature  and  percentage  by  weight 
of  each  element  entering  into  the  composition  of  a  pure  substance. 
It  is  possible,  by  well-known  methods,  to  do  this  for  all  pure 
substances.  As  a  result  of  countless  thousands  of  such  quantita- 
tive experiments  made  by  chemists,  the  conclusion  has  been 
reached  that  the  percentage  composition  of  every  pure  substance 
is  perfectly  definite  for  that  substance  and  is  found  to  be  the  same 
by  whatever  method  we  may  make  the  determination.  This  is 
one  of  the  most  important  laws  of  chemistry.  It  is  usually 
spoken  of  as  the  Law  of  Definite  Composition  or  of  Definite 
Proportions.  This  explains  why  a  pure  substance  always  has 
definite  properties,  from  whatever  source  it  may  be  obtained. 

47o  Hydrogen  and  Its  Gaseous  Compounds. — We  have 
already  become  acquainted  with  hydrogen  and  one  of  its  gaseous 
compounds  hydrogen  chloride,  a  water  solution  of  which  is 
known  as  hydrochloric  acid.  Hydrogen  forms  many  compounds 
which  are  gaseous  at  ordinary  temperatures.  We  shall  now  take 
up  a  study  of  some  of  these,  with  the  object  in  view,  first,  of  dis- 
covering the  nature  of  the  other  element  combined  with  the 
hydrogen;  secondly ?  of  discovering  the  percentage  composition; 
and,  finally,  of  disclosing  a  very  remarkable  relation  between  the 
weights  of  hydrogen  contained  in  equal  volumes  of  these  gases. 

48.  Hydrogen  Chloride. — We  have  found  that  equal  volumes 
of  hydrogen  and  chlorine  combined  to  form  hydrogen  chloride 
gas.  Since  we  know  that  i  liter  of  hydrogen  weighs  o.ogog. 
and  that  i  liter  of  chlorine  weighs  3. 220  g.,  we  find  by  calcula- 
tion that  hydrogen  chloride  contains  2.76  per  cent  by  weight  of 
hydrogen.  By  direct  weighing  of  pure  hydrogen  chloride  gas 
it  is  found  that  i  liter  weighs  i .  642  g.  Since  2 . 76  per  cent  of 
1.642  g.  is  0.045  g->  it  follows  that  i  liter  of  hydrogen  chloride 


34  Introduction  to  General  Chemistry 

gas  contains  o.o45g.  °f  combined  hydrogen.  It  has  already 
been  stated  that  i  liter  of  hydrogen  gas  weighs  0.090  g.,  which 
weight  we  see  is  exactly  double  the  weight  of  hydrogen  in  i  liter  of 
hydrogen  chloride  gas. 

49.  Acetylene:  a  Compound  of  Carbon  and  Hydrogen. — 
Let  us  next  consider  the  gas  acetylene  which  is  extensively  used 
for  illumination.  This  gas  is  obtained  by  allowing  water  to 
drop  on  calcium  carbide.  We  find  that  it  is  a  colorless  gas  with 
a  peculiar  odor.  Everyone  knows  that  it  burns  in  air,  giving  an 
exceedingly  bright  flame.  If  we  collect  and  test  the  products 
coming  from  the  acetylene  flame  we  find  carbon  dioxide  and 
water.  We  find  the  same  products  and  no  others  when  acetylene 
is  burned  in  pure  oxygen  gas,  and  therefore  conclude  that  carbon 


FIG.  22 

and  hydrogen  are  constituents  of  acetylene;  but  the  experiment 
obviously  does  not  decide  whether  oxygen  is  or  is  not  also  a 
constituent  of  acetylene.  This  question  could  be  decided  if  we 
knew  the  percentages  of  carbon  and  hydrogen  in  the  gas. 

50.  The  Analysis  of  Acetylene. — We  may  find  the  per- 
centages of  carbon  and  hydrogen  by  means  of  the  following 
experiment.  A  tube  of  hard  glass  a  centimeter  or  more  in 
diameter  and  30  cm.  long  (Fig.  22)  is  partly  filled  with  pure  dry 
copper  oxide.  The  tube  is  then  heated  red  hot  and  a  measured 
volume  of  acetylene  at  a  known  temperature  and  pressure  is 
caused  to  pass  through  the  tube  and  over  the  heated  copper 
oxide.  It  is  found  that  carbon  dioxide  and  water  are  formed 
and  that  part  of  the  copper  oxide  is  changed  into  metallic  copper. 
A  U-tube  filled  with  calcium  chloride,  for  the  absorption  of  the 
water  formed,  is  attached  to  the  exit  of  the  hard  glass  tube. 
Beyond  this,  attached  by  rubber  tubing,  we  have  a  bulb  contain- 
ing caustic  potash  solution  to  absorb  the  carbon  dioxide.  After 


The  Law  of  Definite  Composition  35 

all  of  the  acetylene  has  been  driven  over  into  the  combustion 
tube  holding  the  copper  oxide,  by  allowing  mercury  from  the 
attached  reservoir  slowly  to  displace  the  acetylene,  a  slow  stream 
of  pure  dry  oxygen  is  passed  into  the  combustion  tube  to  insure 
the  complete  burning  of  the  carbon  of  the  acetylene.  Finally, 
the  oxygen  is  displaced  by  a  stream  of  air. 

The  increase  in  weight  of  the  calcium  chloride  tube  represents 
the  weight  of  water  formed.  Similarly  the  increase  in  weight 
of  the  caustic  potash  bulb  represents  the  weight  of  carbon  dioxide 
obtained.  Now  we  know  that  water  contains  11.2  per  cent  of 
hydrogen  and  that  carbon  dioxide  contains  27.3  per  cent  of 
carbon.  We  may  then  calculate  the  weights  of  hydrogen  and 
carbon  corresponding  to  the  weights  of  water  and  carbon  dioxide 
obtained.  If  we  know  that  i  liter  of  acetylene  under  standard 
conditions,  that  is,  at  o°  and  76  cm.  P,  weighs  i .  190  g.,  we  have 
all  the  data  needed  to  enable  us  to  calculate  the  percentages  of 
hydrogen  and  carbon  in  acetylene.  In  an  actual  lecture  experi- 
ment 200  c.c.  of  pure  dry  acetylene  at  18°  and  75. 4  cm.  gave 
o.i5og.  of  water  and  o.75ig.  of  carbon  dioxide.  From  the 
data  above  we  find  that  the  weight  of  the  acetylene  taken  was 
o.  222  g.,  and  that  the  weights  of  hydrogen  and  carbon  contained 
in  the  water  and  carbon  dioxide  respectively  were  0.0168  g.  and 
0.205  £•>  respectively.  Therefore  acetylene  contains  (according 
to  this  analysis)  7.5  per  cent  of  hydrogen  and  92.3  per  cent  of 
carbon.  The  correct  percentages  are  7 . 7  and  92 . 3  respectively; 
and  since  the  sum  of  these  percentages  is  100,  we  know  that 
hydrogen  and  carbon  are  the  only  elements  contained  in  acetylene. 
We  may  also  calculate  from  the  same  data  the  weight  of  com- 
bined hydrogen  in  one  liter  of  acetylene  under  standard  condi- 
tions: We  find  in  this  way  o .  090  g.  of  hydrogen. 

51.  Ammonia. — Let  us  next  take  up  the  study  of  ammonia. 
Common  household  ammonia,  which  is  familiar  to  everyone,  is  a 
solution  in  water  of  the  substance,  ammonia,  which  is  a  gas  at 
ordinary  temperature  and  pressure.  If  we  warm  such  a  solution 
of  ammonia,  a  gas  having  an  intense  odor  is  given  off.  When 
this  gas,  ammonia,  is  strongly  compressed,  it  condenses  to  a 
colorless  liquid  which  we  speak  of  as  liquid  ammonia.  This  is 


36  Introduction  to  General  Chemistry 

a  commercial  article  which  is  shipped  in  heavy  steel  cylinders  six 
feet  long  and  a  foot  in  diameter.  The  liquid  ammonia  exists 
under  considerable  pressure  in  such  cylinders.  If  the  valve  of 
the  cylinder  is  opened  gaseous  ammonia  escapes.  We  may  use 
a  small  cylinder  of  liquid  ammonia  as  a  convenient  source  of 
ammonia  gas. 

If  we  fill  a  glass  cylinder  with  mercury,  invert  it  in  a  dish  of 
mercury,  and  allow  ammonia  gas  to  escape  under  the  mouth  of 
the  cylinder,  the  mercury  is  displaced  by  the  ammonia  gas.  We 
notice  that  the  gas  is  invisible,  like  air.  It  is  to  be  distinguished 
from  air,  however,  by  its  intense  odor,  as  well  as  in  other  ways. 
If  we  dip  the  mouth  of  the  cylinder,  which  has  been  closed  by  a 
glass  plate,  into  a  vessel  of  water,  we  find  that  the  water  rushes 
into  the  cylinder  almost  as  readily  as  if  the  space  were  a  vacuum. 
An  examination  of  the  water  now  shows  that  it  has  new  prop- 
ertiess  The  water  now  has  the  odor  of  ammonia,  it  has  a 
peculiar  disagreeable  taste,  and  changes  the  color  of  immersed 
red  litmus  paper  blue.  If  we  bring  a  burning  candle  into  a 
cylinder  of  ammonia  the  flame  of  the  candle  is  extinguished  but 
the  ammonia  does  not  take  fire.  These  properties  distinguish 
ammonia  from  oxygen,  hydrogen,  and  acetylene. 

52.  Ammonia  a  Compound  of  Nitrogen  and  Hydrogen. — We 
may  now  inquire,  What  is  the  chemical  composition  of  ammonia? 
Is  it  an  elementary  substance  or  a  compound,  and,  if  a  compound, 
of  what  elements  is  it  composed?  If  ammonia  gas  is  passed 
through  a  heated  glass  tube  containing  copper  oxide  we  observe 
that  a  colorless  liquid  condenses  in  the  cold  part  of  the  tube. 
This  liquid  proves  to  be  water.  We  find  also  that  a  colorless, 
odorless  gas  is  formed.  If  we  pass  this  gas  into  limewater  we 
observe  no  result  and  conclude,  therefore,  that  this  gas  is  not 
carbon  dioxide.  We  find  that  the  gas  is  not  appreciably  soluble 
in  water,  so  that  it  cannot  be  unchanged  ammonia  gas.  If  we 
test  the  gas  with  a  burning  candle  we  find  that  it  neither  burns 
nor  supports  combustion.  The  student  will  doubtless  recall 
(10)  that  this  gas  has  just  those  properties  which  the  portion 
of  the  air  left  after  the  removal  of  oxygen  by  mercury  or  phos- 
phorus possesses.  It  would  seem,  therefore,  to  be  nitrogen. 


The  Law  of  Definite  Composition  37 

The  identity  of  the  gas  with  nitrogen  is  confirmed  by  a  deter- 
mination of  the  density;  whereupon  it  is  found  that  a  liter 
weighs  i.25ig.  Since  water  and  copper  were  formed  from 
ammonia  and  copper  oxide,  we  conclude  that  ammonia  has 
furnished  the  hydrogen  which  united  with  the  oxygen  supplied 
by  the  copper  oxide  to  form  the  water  obtained  in  the  preceding 
experiment.  Ammonia  must  be  a  compound  containing  nitrogen 
and  hydrogen.  It  has  been  shown  in  many  ways  by  experiments, 
which  we  need  not  consider  at  present,  that  nitrogen  and  hydro- 
gen are  the  only  constituents  of  ammonia. 

53.  The  Percentage  Composition  of  Ammonia. — The  per- 
centage of  hydrogen  in  ammonia  may  be  found  by  carrying 
out  the  experiment  above  described  with  a  known  volume  of 
ammonia  measured  at  a  known  temperature  and  pressure.     If 
we  cause  the  ammonia  to  pass  through  the  heated  copper  oxide 
tube,  driving  out  water  vapor  completely  by  means  of  air  after 
all  of  the  ammonia  has  passed  into  the  tube,  and  if  the  products 
are  caused  to  pass  through  a  calcium  chloride  tube  connected 
to  the  copper  oxide  tube  as  in  the  determination  of  the  composi- 
tion of  acetylene,  the  increase  in  weight  of  the  calcium  chloride 
tube  gives  us  the  weight  of  water  formed  from  the  hydrogen  of 
the  ammonia  used.     Knowing  as  we  do  the  percentage  of  hydro- 
gen in  water,  if  we  know  the  weight  of  a  liter  of  ammonia  gas 
(o.772g.)   we  may  calculate  the  percentage  of  hydrogen  in 
ammonia  and  also  the  weight  of  combined  hydrogen  in  i  liter  of 
ammonia  gas  measured  under  standard  conditions.    We  find 
this  latter  weight  to  be  o.  135  g. 

54.  Methane,  Another  Compound  of  Carbon  and  Hydro- 
gen.— The  chief  component  of  natural  gas  is  a  substance  called 
methane.    This  same  gas  methane  often  escapes  in  bubbles  when 
the  decaying  vegetable  matter  in  marshes  is  disturbed.     For  this 
reason  methane  is  also  known  as  marsh  gas.    We  may  prepare 
methane  artificially  in  the  laboratory  by  methods  which  we  need 
not  now  discuss.    It  may  be  collected  over  water,  as  its  solu- 
bility in  water  is  slight.    We  note  that  it  is  a  colorless  gas,  that 
it  is  lighter  than  air,  since  the  gas  will  escape  rapidly  from  an 
open  cylinder  when  the  mouth  of  the  cylinder  is  turned  upward, 


38  Introduction  to  General  Chemistry 

but  will  not  escape  if  the  mouth  is  downward.  One  liter  of 
methane  weighs  0.721  g.,  which  is  but  little  more  than  half  of 
the  weight  of  the  same  volume  of  air.  If  we  bring  a  lighted 
candle  into  a  cylinder  of  methane  we  find  that  the  gas  burns 
with  a  slightly  luminous  flame  but  that  the  candle  flame  is 
extinguished. 

55.  The  Quantitative  Analysis  of  Methane. — If  we  examine 
the  products  of  combustion  from  a  methane  flame  we  find  .water 
and  carbon  dioxide,  from  which  we  know  that  methane  is  a  com- 
pound of  carbon  and  hydrogen  with  or  without  other  elements. 
We  may  determine  the  quantitative  composition  of  methane  by 
precisely  the  same  method  as  that  used  for  the  quantitative 
analysis  of  acetylene,  whereupon  we  find  that  methane  contains 
75.0  per  cent  of  carbon  and  25.0  per  cent  of  hydrogen  by  weight. 
Since  the  sum  of  these  percentages  is  100  we  know  that  methane 
must  contain  only  the  elements  carbon  and  hydrogen.    From  the 
data  obtained  in  the  analysis  of  methane  we  may  also  calculate 
that  i  liter  of  methane  under  standard  conditions  contains 
o.  1 80  g.  of  combined  hydrogen. 

56.  The  Weight  of  Hydrogen  in  One  Liter  of  Gaseous  Hydro- 
gen Compounds. — By  a  study  of  the  composition  of  the  four 
gases,  hydrogen  chloride,  acetylene,  ammonia,  and  methane,  as 
well  as  of  hydrogen  itself,  we  have  found  the  weight  of  hydrogen 
in  i  liter  of  each.    These  results  may  now  be  tabulated  as  in 
Table  II.     An  inspection  of  the  results  given  in  the  table  reveals 

TABLE  II 

Hydrogen  chloride o .  045  g. 

Hydrogen o .  090 

Acetylene o .  090 

Ammonia o.  135 

Methane 0.180 

a  remarkable  fact.  The  weight  of  hydrogen  in  i  liter  of  hydrogen 
chloride  is  less  than  that  in  any  other  case.  The  weight  per  liter 
of  hydrogen  gas  itself  is  double  the  weight  of  hydrogen  in  i  liter  of 
hydrogen  chloride.  Likewise  the  weight  of  hydrogen  in  i  liter  of 
acetylene  is  exactly  equal  to  the  weight  of  a  liter  of  free  hydrogen  and 


The  Law  of  Definite  Composition  39 

also  double  the  weight  of  hydrogen  in  i  liter  of  hydrogen  chloride. 
The  weight  of  hydrogen  in  i  liter  of  ammonia  is  three  times  that  in 
i  liter  of  hydrogen  chloride,  while  in  the  case  of  methane  the  weight 
of  hydrogen  per  liter  is  four  times  the  weight  of  this  element  in  the 
same  volume  of  hydrogen  chloride. 

If  we  consider  the  weight  of  hydrogen  in  a  liter  of  hydrogen 
chloride  as  unity,  we  find  that  the  weights  in  the  same  volumes  of  the 
other  gases  are  expressed  by  the  numbers  2,  3,  or  4.  It  is  obvious 
that  the  relations  we  discussed  would  also  hold  equally  well  if 
we  dealt  with  weights  of  hydrogen  contained  in  any  other  fixed 
volume,  as  a  cubic  foot  or  a  cubic  meter.  We  could  express  the 
facts  by  saying  that  the  weight  of  hydrogen  contained  in  a  fixed 
volume  of  any  of  these  gases  is  in  each  case  a  multiple  of  the  mini- 
mum weight,  which  is  found  in  the  case  of  hydrogen  chloride  gas. 
Since  i  liter  of  hydrogen  chloride  gas  contains  0.045  £•  °f  hydro- 
gen, i  g.  of  combined  hydrogen  would  be  contained  in  22.4 
liters1  of  hydrogen  chloride.  In  the  same  volume  of  the  other  gases 
the  weights  of  hydrogen  would  be  2  g.,  3  g.,  or  4  g. 

1  In  reality  1-7-0.045  gives  22. 2  instead  of  the  correct  value  22.4  liters.  The 
discrepancy  is  caused  by  the  fact  that  the  numbers  used  are  only  approximate. 
This  subject  is  discussed  further  in  section  222. 


CHAPTER  V 
SYMBOLS  AND  CHEMICAL  FORMULAE 

57.  Gaseous  Carbon  Compounds. — We  may  now  inquire 
whether  the  remarkable  relations  between  the  weights  of  hydro- 
gen in  equal  volumes  of  compounds  of  hydrogen  hold  good  in  the 
case  of  compounds  of  other  elements.    We  have  already  studied 
three  gaseous  compounds  of  carbon:   carbon  dioxide,  acetylene, 
and  methane,  and  have  seen  how   the  percentage   composi- 
tion of  each  is  determined.     Before  discussing  the  results  so 
obtained,  let  us  consider  two  new  gaseous  compounds  of  carbon : 
propane  and  trimethylamine. 

58.  Propane:    a   Compound   of   Carbon  and   Hydrogen.— 
Propane  is  found  in  small  amounts  in  the  natural  gas  of  some 
wells  and  also  dissolved,  in  small  quantities,  in  crude  petroleum. 
It  may  also  be  obtained  artificially  by  methods  well  known  to  the 
chemist,  the  nature  of  which  we  need  not  now  consider.     We 
observe  that  propane  is  a  colorless,  odorless  gas  which  is  some- 
what heavier  than  air,  i  liter  under  standard  conditions  weighing 
i .  97  g.     We  find  that  propane  resembles  methane  in  its  chemi- 
cal behavior,  since  it  extinguishes  a  burning  candle  but  takes 
fire  itself  at  the  same  time,  burning  with  a  slightly  luminous 
flame  and  yielding  carbon  dioxide  and  water  as  the  only  products 
of  combustion.     The  analysis  of  propane  may  be  carried  out  in 
precisely  the  same  manner  as  our  analysis  of  methane  and  acety- 
lene.    We  find  in  this  way  that  propane  contains  81 .8  per  cent 
of  carbon  and  18.2  per  cent  of  hydrogen.     Since  the  sum  of  these 
percentages  is  100,  it  follows  that  carbon  and  hydrogen  are  the 
only  constituents  of  propane. 

59.  Trimethylamine:  a  Compound  of  Carbon,  Hydrogen,  and 
Nitrogen. — Trimethylamine  is  a  colorless  gas  about  twice  as 
heavy  as  air,  i  liter  weighing  2 . 65  g.     Its  odor  is  very  powerful 
and  somewhat  disagreeable,  but  if  inhaled  in  small  quantities  the 
gas  is  not  poisonous  nor  irritating,  as  is,  for  example,  chlorine  gas. 
The  odor  is  that  of  decaying  fish.     In  fact,  the  gas  can  be 

40 


Symbols  and  Chemical  Formulae 


obtained  from  products  separated  from  herring  brine.  We  find 
that  the  gas  is  very  easily  soluble  in  water  and  that  the  solution 
turns  red  litmus  paper  blue,  just  as  ammonia  does;  but  the  gas 
may  be  distinguished  from  ammonia  by  the  fact  that  it  will  burn, 
whereas  ammonia  will  not.  It  is  easy  to  discover  that  water  and 
carbon  dioxide  are  formed  when  trimethylamine  is  burned  in  air 
or  in  oxygen.  If  we  pass  trimethylamine  through  a  tube  con- 
taining heated  copper  oxide  we  obtain,  in  addition  to  water  and 
carbon  dioxide,  a  colorless,  odorless,  incombustible  gas  which  can 
easily  be  identified  as  nitrogen.  These  facts  show  that  tri- 
methylamine contains  the  elements  carbon,  hydrogen,  and  nitrogen. 
We  could  determine  the  percentages  of  carbon  and  hydrogen  by 
finding  the  weights  of  carbon  dioxide  and  water  formed  by  the 
action  of  the  gas  on  hot  copper  oxide,  as  in  analyses  previously 
made.  We  might  also  find  the  percentage  of  nitrogen  by  finding 
the  volume  of  nitrogen  which  we  could  obtain  from  a  known 
volume  of  the  gas.  The  percentages  of  carbon,  hydrogen,  and 
nitrogen  would  be  found  to  be  61  .o,  15.3,  and  23 . 7  respectively. 
60.  The  Weights  of  Carbon  in  i  Liter  and  in  22 . 4  Liters  of 
Gaseous  Carbon  Compounds. — Let  us  now  consider  the  facts 
presented  in  Table  III.  The  weight  of  i  liter  and  the  percentage 

TABLE  III 


Weight  of 
i  Liter 

Percentage  of 
Carbon 

Weight  of 
Carbon  in 
i  Liter 

Weight  of 
Carbon  in 
22  .  4  Liters 

Methane 

O   72 

7  r    Q 

O    ^4. 

1  2 

Carbon  dioxide  
Acetylene 

I.Q7 
I    IQ 

27-3 
02    ^ 

0-54 
I    08 

12 
24 

Propane 

I    07 

81  8 

I    62 

l6 

Trimethylamine  

2    6< 

61  o 

I    62 

*6 

of  carbon  in  each  of  the  five  gaseous  compounds  of  carbon  we 
have  studied  are  given  in  the  first  and  second  columns  of  figures. 
The  product  of  the  weight  of  i  liter  of  a  gas  by  the  percentage  of 
carbon  it  contains  gives  the  weight  of  combined  carbon  in  i  liter. 
These  products  are  given  in  the  third  column.  The  weights  of 
carbon  in  22.4  liters,  as  given  in  the  last  column,  are  found  by 
multiplying  the  corresponding  weights  in  the  third  column  by 
22.4. 


42  Introduction  to  General  Chemistry 

We  see  by  a  glance  at  the  last  column  of  the  table  that  22.4 
liters  of  carbon  dioxide  and  methane  contain  12  g.  of  combined 
carbon,  that  the  same  volume  of  acetylene  contains  24  g.  of 
carbon,  while  the  weight  of  combined  carbon  in  22.4  liters  of 
propane  and  trimethylamine  is  36  g.,  and  therefore  that  the 
•weight  of  carbon  in  22 . 4  liters  of  any  of  these  gases  is  either  one, 
two,  or  three  times  12  g.  In  the  case  of  gaseous  hydrogen  com- 
pounds, we  found  that  the  weight  of  hydrogen  was  either  one, 
two,  three,  or  four  times  i  g.,  which  was  the  minimum  weight 
of  this  element  found  in  any  case.  We  thus  find  that  in  22 . 4 
liters  of  various  pure  gases  the  minimum  weight  of  hydrogen  is  i  g. 
and  the  minimum  weight  of  carbon  12  g.,  and,  further,  that  if  a 
greater  weight  of  either  of  these  elements  is  contained  in  this  volume 
of  any  pure  gas}  the  weight  is  a  multiple  of  the  minimum  weight 
by  a  small  whole  number. 

Let  us  now  consider  the  weights  of  carbon  and  hydrogen 
contained  in  22.4  liters  of  the  three  gaseous  compounds  which 
contain  only  carbon  and  hydrogen,  namely,  methane,  acetylene, 
and  propane.  In  22.4  liters  of  methane  we  find  12  g.  of  carbon 
combined  with  4  g.  of  hydrogen.  In  the  same  volume  of  acety- 
lene, 24  g.  of  carbon  combined  with  2  g.  of  hydrogen,  and  in  the 
case  of  propane  36  g.  of  carbon  combined  with  8  g.  of  hydrogen. 
Without  considering  at  present  the  theoretical  significance  of 
the  remarkable  facts  which  these  figures  show,  we  may  consider 
a  practical  application  of  the  facts  which  will  enable  us  to  express 
the  composition  of  these  gases  in  a  simple  fashion. 

The  student  must  realize  that  since  we  have  three  compounds 
all  consisting  of  carbon  and  hydrogen  and  having  different  prop-, 
erties,  the  difference  in  percentage  composition  must  be  an 
important  factor  in  determining  the  properties  of  the  substance. 
He  will  also  understand  that  a  knowledge  of  the  percentage  com- 
position is  a  matter  of  prime  importance  for  the  chemist,  and  that 
any  scheme  by  means  of  which  a  knowledge  of  the  composition 
by  weight  could  be  easily  memorized  would  be  important. 

61.  Symbols. — Suppose  we  represent  i  g.  of  hydrogen  by  a 
sign  or  symbol  and  choose  the  letter  H  for  this  purpose.  We 
could,  then,  represent  by  H  taken  four  times  the  weight  of  hydro- 


Symbols  and  Chemical  Formulae  43 

gen  contained  in  22.4  liters  of  methane;  by  H  taken  twice,  or 
2H,  the  amount  of  hydrogen  in  22.4  liters  of  acetylene;  and 
similarly  by  8H,  the  amount  of  hydrogen  in  22 . 4  liters  of  propane. 
Suppose  that,  on  the  other  hand,  we  represent  12  g.  of  carbon  by 
the  sign  or  symbol  C,  then  C,  2C,  and  30  will  represent  the 
weights  of  carbon  in  22.4  liters  of  methane,  acetylene,  and 
propane  respectively.  The  weights  of  carbon  and  hydrogen  in 
22.4  liters  of  methane  may  then  be  represented  by  writing  iC 
together  with  qH.  As  a  matter  of  convenience  the  multiples  i 
for  the  C  and  4  for  the  H,  are  written  as  subscripts;  so  that 
instead  of  iC  and  4!!  we  write  CiH4.  In  practice  no  subscript 
is  used  when  the  multiple  is  i.  The  composition  of  methane  is 
represented  simply  by  CH4. 

62.  Chemical  Formulae. — In  a  similar  way  we  may  represent 
the  weights  of  carbon  and  hydrogen  in  22 . 4  liters  of  acetylene  by 
C2H2  while  the  composition  of  the  same  volume  of  propane  may 
be  represented  by  C3H8.  We  call  H  the  symbol  for  hydrogen,  and 
for  the  present  we  may  consider  that  H  or  iH  represents  i  g. 
of  hydrogen  and  similarly  that  C,  the  symbol  for  carbon,  repre- 
sents 12  g.  of  that  element.  We  call  the  expressions  CH4,  C2H2, 
and  C3Hg  the  formulae  of  methane,  acetylene,  and  propane 
respectively.  We  shall  now  proceed  to  show  how  this  system 
may  be  extended  to  all  gaseous  compounds  of  any  element  what- 
ever. 

Chemists  are  familiar  with  a  large  number  of  gases  in  addi- 
tion to  those  which  we  have  already  studied.  Some  of  these  are 
of  much  practical  importance  while  others  are  chiefly  of  interest 
to  the  chemist  for  scientific  reasons.  In  every  case  it  is  a  simple 
matter  to  determine  the  weight  of  i  liter  of  the  gas  under  stand- 
ard conditions,  the  method  of  making  the  determination  being 
essentially  the  same  in  all  cases.  Furthermore,  by  methods 
which  are  well  known  to  chemists  we  may  determine  what  ele- 
ments compose  any  gas,  and  by  means  of  a  quantitative  analysis 
we  may  determine  the  percentage  of  each  element  in  the  gas. 
If  we  calculate  in  the  case  of  each  gas  the  weight  of  each  element 
contained  in  22.4  liters  of  the  gas,  we  obtain  results  like  those 
shown  in  Table  IV. 


44 


Introduction  to  General  Chemistry 


63.  The  Minimum  \Veights  of  Oxygen,  Nitrogen,  and 
Chlorine. — An  inspection  of  the  results  given  in  Table  IV  shows 
that  the  same  regularity  in  the  weights  of  hydrogen  and  carbon 
holds  in  all  cases,  as  we  have  observed  it  to  hold  in  the  few  cases 
discussed  in  the  preceding  paragraphs.  We  notice  also  that  the 
minimum  weight  of  oxygen  in  22.4  liters  of  any  of  its  gaseous 

TABLE  IV 

WEIGHTS  OF  CONSTITUENTS  IN  22.4  LITERS  OF  GASES 


Substance 

Oxygen 

Hydrogen 

Carbon 

Nitrogen 

Chlorine 

Formula 

Oxygen  .  .  . 

2X16 

O2 

Carbon  monoxide    .... 

1X16 

1X12 

CO 

Carbon  dioxide  

2X16 

lXl2 

CO2 

Nitrous  oxide  

1X16 

2X14 

N2O 

Nitric  oxide  

1X16 

1X14 

NO 

Nitrosyl  chloride  
Hypochlorous  oxide  
Chlorine  dioxide 

1X16 
1X16 
2X16 

1X14 

1X35.5 
2X35.5 
iX^s    s 

NOC1 
C120 
C1O2 

Phosgene 

1X16 

lXl2 

2X^  <; 

COC12 

Methyl  ether  .    .  . 

1X16 

6X1 

2Xl2 

C2H6O 

Hydrogen  .... 

2X1 

H2 

Hydrogen  chloride.  . 

1X1 

iX^  * 

HC1 

Prussic  acid  
Ammonia  
Methane  
Acetylene  
Ethylene  
Ethane  

1X1 
3X1 
4X1 
2X1 
4X1 
6X1 

IXI2 

IXI2 
2X12 
2X12 
2Xl2 

1X14 
1X14 

HNC 
NH3 
CH, 
C2H2 
C2H4 
C2H6 

Propylene  
Propane 

6X1 
8X1 

3X12 
3X12 

C3H6 
CjHs 

Methyl  chloride" 

3X1 

lXl2 

iX^<?  ^ 

CH3C1 

Ethyl  chloride 

^Xi 

2Xl2 

iX^J  «? 

C2HSC1 

Methylamine 

^Xi 

lXl2 

1X14 

CHSN 

Nitrogen 

2X14 

N2 

Cyanogen 

2Xl2 

2X14 

C2N2 

Cyanogen  chloride.  . 

lXl2 

1X14 

iX^."? 

C1NC 

Chlorine      .  .  . 

2X35.5 

C12 

Trimethylamine  

9X1 

3X12 

1X14 

c3Eyff 

compounds  is  16  g.,  and  that  this  weight  is  found  in  many 
cases,  while  in  others  the  weight  is  twice  16.  In  the  case  oi 
the  compounds  of  nitrogen  we  note  that  the  minimum  weight 
is  14  g.  and  that  in  other  cases  the  weight  is  double  this  mini- 
mum weight.  In  the  case  of  chlorine  compounds  the  minimum 
weight  of  chlorine  is  35.5  g.,  while  those  compounds  with 
a  larger  proportion  of  chlorine  contain  double  the  minimum 
weight. 


Symbols  and  Chemical  Formulae  45 

64.  The  Law  of  Minimum  and  Multiple  Weights. — Entirely 
analogous  regularities  will  be  found  if  we  consider  the  data 
obtained  from  a  study  of  the  gaseous  compounds  of  any  other 
elements.     For  each  element  we  find  a  minimum  weight  in  the 
volume  0/22.4  liters  of.  any  of  its  gaseous  compounds  under  standard 
conditions  and  also  find  that  the  weight  if  greater  than  the  minimum 
would  be  2,  3,  or  4,  or  some  small  multiple  of  this  minimum.    This 
last  statement  may  be  called  the  Law  of  Minimum  and  Multiple 
Weights. 

65.  The  Chemical  Unit  Volume:    22.4  Liters. — The  volume 
22.4  liters  thus  becomes  a  kind  of  unit  volume  for  the  chemist, 
this  particular  volume  having  been  chosen  because  it  contains 
i  g.  of  hydrogen  in  the  case  of  those  hydrogen  compounds 
which  contain  the  minimum  weight  of  this  element.     In  this 
volume  no  other  element  has  a  minimum  weight  as  small  as  that 
of  hydrogen. 

66.  Symbols  Represent  Minimum  Weights. — In  the  same 
manner  as  that  suggested  in  a  preceding  paragraph  for  hydrogen 
and  carbon,  we  may  represent  the  minimum  weight  of  each  of 
the  other  elements  by  a  symbol.    Table  V  shows  the  minimum 
weights  of  the  five  elements  we  have  been  considering,  together 
with  the  corresponding  symbols. 

TABLE  V 
MINIMUM  WEIGHTS  IN  22.4  LITERS,  AND  SYMBOLS 


Hydrogen  

I.Og. 

H 

Carbon  

12.0 

C 

Nitrogen  

14.0 

N 

Oxygen 

16  o 

o 

Chlorine 

•2C    c 

Cl  ~~ 

67.  Making  Formulae. — We  see  from  Table  IV  that  22.4 
liters  of  carbon  dioxide  contain  1 2  g.  of  carbon  combined  with 
2X16  g.  of  oxygen.  We  may,  therefore,  represent  the  composi- 
tion of  the  quantity  of  carbon  dioxide  in  22.4  liters  by  the 
formula  C02.  In  an  analogous  fashion  we  may  obtain  as  the 
formula  representing  the  composition  of  22.4  liters  of  ammonia, 
NH3,  and  as  the  formula  for  hydrogen  chloride,  HC1.  By  making 


46  Introduction  to  General  Chemistry 

use  of  this  system  the  student  will  now  have  no  difficulty  in 
writing  down  at  once  the  formula  of  each  of  the  gases  from  the 
data  contained  in  Table  IV.  He  will  also  readily  see  that  it  is  a 
much  less  difficult  task  to  learn  the  formulae  of  such  gases  than 
to  learn  their  percentage  composition;  that  is  to  say,  it  is  an  easier 
tax  upon  the  mind  to  remember  the  formula  HC1  than  to  remem- 
ber that  hydrogen  chloride  contains  2.76  per  cent  of  hydrogen 
and  97 . 24  per  cent  of  chlorine. 

68.  The  Practical  Use  of  Formulae. — A  review  of  the  methods 
employed  in  arriving  at  the  results  represented  by  the  formula 
of  any  substance  shows  that  in  each  case  we  have  made  use  of  the 
knowledge,  first,  of  the  weight  of  i  liter  of  the  gas,  and,  secondly, 
of  the  percentage  of  each  of  its  elementary  constituents.     Con- 
versely, if  we  know  the  weights  which  the  symbols  of  the  elements 
represent,  and  know  the  formula  of  a  gas,  we  may  by  working 
backward  find  its  percentage  composition.     For  example,  sup- 
pose that  we  remember  that  the  formula  of  methane  is  CH4 
and  know  that  H  stands  for  i  g.  of  hydrogen  and  C  for  1 2  g. 
of  carbon.    Then  22.4  liters  of  methane  contain  4 g.  hydrogen 
combined  with  1 2  g.  of  carbon.    The  proportion  of  hydrogen  is, 
therefore,  4/16,  or  25  per  cent,  and  of  carbon  12/16,  or  75  per  cent, 
the  weight  of  22 . 4  liters  being  16  g.,  and  i  liter  weighs  16/22 . 4= 
0.72  g.     In  calculating  in  this  way  the  density  and  percentage 
composition  of  methane  we  are  merely  reproducing  the  results 
which  originally  were  obtained  by  experiment.    In  order  to  find 
the  formula  of  any  gas,  we  must  know  its  density  and  the  per- 
centage of  each  elementary  constituent.    We  find  by  actual 
experience  that  we  can  represent  by  a  formula,  usually  of  a  very 
simple  character,  the  composition  of  22.4  liters  of  any  gaseous 
substance. 

69.  Formulae  of  Liquids  and  Solids. — The  system  which  we 
have  just  considered  is  capable  of  extension  to  liquid  and  solid 
substances,  in  which  case,  however,  the  formula  may  have  a 
slightly  less  definite  meaning.    We  may  illustrate  this  by  con- 
sidering the  cases  of  water  and  mercury  oxide.    We  have  found 
that  water  is  composed  of  11.2  per  cent  of  hydrogen  and  88.8 
per  cent  of  oxygen,  from  which  we  observe  that  the  weight  of 


Symbols  and  Chemical  Formulae  47 

the  oxygen  is  8  times  the  weight  of  the  hydrogen  with  which  it  is 
united.  This  ratio  of  hydrogen  and  oxygen  might  be  represented 
by  H2O,  since  this  formula  would  mean  that  2  g.  of  hydrogen  are 
united  with  16  g.  of  oxygen,  which  weights  of  hydrogen  and 
oxygen  are  in  the  ratio  of  i  to  8,  but  the  formulae  H402  and 
H6O3  would  also  represent  equally  well  the  proportion  of  hydro- 
gen and  oxygen  actually  found  in  water. 

70.  The  Formula  of  Water. — We  may  be  led  to  choose  a  con- 
sistent formula  for  water  by  the  consideration  of  the  density  of 
water  vapor  or  steam;  but  in  this  case  the  density  determination 
must  be  made  at  a  temperature  above  the  boiling-point  of  water, 
if  we  work  at  atmospheric  pressure.     Since  the  effect  of  changes 
of  pressure  and  temperature  upon  the  volume  of  a  given  quantity 
of  steam  are  the  same  as  upon  an  equal  volume  of  any  gas  which 
would  not  liquify  if  cooled  to  o°  at  76  cm.  pressure,  we  might 
calculate  by  the  laws  of  Boyle  and  Charles  what  the  volume  of 
the  known  weight  of  steam  measured  at  a  high  temperature  and 
known  pressure  would  be  if  the  steam  were  under  standard  con- 
ditions, that  is,  at  o°  and  76  cm.  pressure.     It  has  been  found  in 
this  way  that  i  liter  of  water  vapor  if  it  did  not  condense  to  a 
liquid  would  weigh  o.8o6g.  under  standard  conditions,  which 
corresponds  to  a  weight  of  18  g.  for  22.4  liters.     Now,  11.2  per 
cent  of  18  g.  is  2  g.  and  88.8  per  cent  of  18  g.  is  16  g.     From 
these  results  we  conclude  that  if  water  vapor  could  exist  under 
standard  conditions  as  a  gas  that  22.4  liters  would  contain 
2  g.  of  combined  hydrogen  and  16  g.  of  combined  oxygen,  which 
amounts  would  be  exactly  represented  by  the  formula  H20. 

71.  Formulae  of  Volatile  Liquids  and  Solids. — In  a  perfectly 
analogous  fashion  we  could  find  the  formula  for  any  other  volatile 
substance,  the  density  of  whose  vapor  we  could  measure  experi- 
mentally.   Such  a  procedure  would  enable  us  to  represent  by  a 
formula  the  composition  of  a  great  number  of  volatile  chemical 
substances  which  are  not  gaseous,  but  are  liquid  or  solid  under 
ordinary  conditions  of  temperature  and  pressure. 

72.  Formulae  of  Involatile  Substances. — There  are,  however, 
many  chemical  substances  which  are  not  volatile  or  which  can- 
not be  volatilized  at  temperatures  at  which  we  could  make 


48  Introduction  to  General  Chemistry 

experimental  determinations  of  their  vapor  densities.  There 
are  other  solids  and  liquids  which  would  be  decomposed  if 
strongly  heated.  For  such  substances  we  could  not  find  chemi- 
cal formulae  in  the  same  way  as  for  gases  or  volatile  substances. 
However,  we  can  and  do  represent  by  formulae  the  composition 
of  such  involatile  substances. 

73.  The  Formula  of  Red  Oxide  of  Mercury.— The  method  of 
obtaining  the  formula  of  such  a  substance  may  be  illustrated  by 
the  case  of  the  red  oxide  of  mercury,  which,  it  will  be  remem- 
bered, is  readily  decomposed  when  heated  into  mercury  and 
oxygen.    We  find  by  analysis  that  this  compound  contains 
92.6  per  cent  of  mercury  and  7.4  per  cent  of  oxygen.     By  the 
experimental  study  of  volatile  mercury  compounds,  as  well  as  of 
mercury  itself,  we  find  that  the  minimum  weight  of  mercury  in 
22.4  liters  is  200  g.,  and  therefore  represent  this  weight  of 
mercury  by  the  symbol  Hg.    It  now  remains  to  discover  what 
multiples  of  200  for  the  mercury  and  of  16  for  the  oxygen  are 
in  the  same  ratio  as  the  percentages  of  mercury  and  oxygen  in 
mercury  oxide.    We  find  very  easily  that  200  is  to  1 6  as  92.6  is 
to  7 . 4,  and  from  this  we  write  the  formula  HgO. 

We  could  of  course  represent  the  same  proportions  of  mercury 
and  oxygen  by  the  formula  Hg2O2.  But  we  are  not  able  to 
decide  which  of  these  to  choose  as  in  the  case  of  a  volatile  sub- 
stance where  the  formula  represents  the  quantity  of  material  in 
22.4  liters  of  the  gas  or  vapor  under  standard  conditions.  In 
such  a  case  we  choose  the  simpler  formula,  in  this  case  HgO,  but 
we  must  bear  in  mind  that  the  formula  does  not  mean  quite  as 
much  in  such  a  case  as  in  that  of  a  gas  or  volatile  substance, 
where  it  always  represents  in  addition  to  the  true  proportion 
of  the  constituent  elements  the  actual  weights  of  each  in  22.4 
liters  of  the  gas  under  standard  conditions. 

74.  Symbol  Weights  and  Formula  Weights.— The  letter  or 
pair  of  letters  which  represents  the  minimum  weight  of  an 
element  in  22.4  liters  of  any  of  its  gaseous  compounds  is  called 
the  symbol  of  that  element  and  the  weight  which  this  symbol 
represents  may  then  be  called  the  symbol  weight.    Each  of  the 
eighty-five  or  more  known  elements  has  been  assigned  a  definite 


Symbols  and  Chemical  Formulae  49 

symbol  which  represents  a  definite  symbol  weight.  We  have 
seen  (62)  how  the  quantities  of  each  element  in  22.4  liters  of  a 
compound  gas  may  be  represented  by  a  formula  made  up  of 
symbols,  each  symbol  being  multiplied  by  a  factor  which  shows 
how  many -times  the  minimum  weight  of  the  element  is  present 
in  22 . 4  liters  of  the  gaseous  compound.  The  sum  of  the  weights 
represented  by  the  various  symbols  each  multiplied  by  its  factor 
is  naturally  the  weight  of  22.4  liters  of  the  gas,  represented  by 
the  formula.  This  weight  is  often  spoken  of  as  the  formula 
weight.  In  the  case  of  an  involatile  solid  substance  the  formula 
weight  is  the  weight  represented  by  the  formula  but  indicates 
only  theoretically  the  weight  which  we  should  expect  22.4  liters 
of  the  substance  to  have  if  it  were  a  gas  under  standard  con- 
ditions. 

75.  The  Formulae  of  Some  Elementary  Gases; — It  is  impor- 
tant to  note  that  22.4  liters  of  the  gases  hydrogen,  oxygen, 
nitrogen,  and  chlorine  weigh  2,  32,  28,  and  71  g.  respectively 
(63,  Table  IV).  These  weights  are  for  each  element  just  double 
the  minimum  weights  which  we  find  in  numerous  compounds 
of  the  elements  and  therefore  in  each  case  just  double  the  weight 
represented  by  the  symbol.  We  must  therefore  write,  as  the 
formulae  of  these  gases,  H2,  O2,  N2,  and  C12,  respectively.  The 
formula  of  an  elementary  gas  in  the  free  state  will  then  represent 
the  quantity  of  that  gas  in  22.4  liters.  We  must  here  point 
out  that  not  every  element  in  the  form  of  gas  or  vapor  is  to  be 
represented  by  a  formula  composed  of  its  symbol  taken  twice. 
For  example,  the  vapors  of  mercury  and  sodium  have  the  single 
symbol  formulae  Hg  and  Na,  respectively;  on  the  other  hand, 
the  formulae  of  the  vapors  of  the  elements  phosphorus  and  sulfur 
are  P4  and  Sg. 


CHAPTER  VI 
CHEMICAL  EQUATIONS 

76.  Equations.  —  In  this  chapter  we  shall  see  how  it  is  possible 
to  represent  in  a  very  simple  way  the  quantities  of  substances 
entering  into  and  formed  in  a  chemical  reaction.  Let  us  con- 
sider the  case  of  hydrogen  and  chlorine  which  has  already  been 
studied  experimentally.  We  have  learned  that  hydrogen  and 
chlorine  unite  to  form  hydrogen  chloride  (44).  Furthermore 
we  find  by  experiment  that  one  volume  of  hydrogen  and  one 
volume  of  chlorine  give  two  volumes  of  hydrogen  chloride;  so 
that  if  22.4  liters  of  hydrogen  united  with  22.4  liters  of  chlorine 
we  should  obtain  44.8  liters  of  hydrogen  chloride.  Now  we 
may  represent  22.4  liters  of  hydrogen  by  the  formula  H2  and 
22.4  liters  of  chlorine  by  C12,  while  for  twice  22.4  liters  of  hydro- 
gen chloride  we  put  the  coefficient,  2,  in  front  of  the  formula  and 
write  2HC1.  We  may  then  express  the  facts  by  stating  that  H2 
plus  C12  gives  2HC1  or 

Ha+Cl3-»2HCl. 

which  may  also  be  written 


We  call  either  of  these  expressions  the  equation  for  the  reaction 
between  hydrogen  and  chlorine. 

77.  What  an  Equation  Means.  —  The  equation 

H2+C12->2HC1 

expresses  the  fact  that  the  quantity  of  hydrogen  represented  by 
the  formula  H2  or  2  g.  unites  with  the  quantity  of  chlorine 
represented  by  C12  or  71  g.  to  give  the.  quantity  of  hydrogen 
chloride  represented  by  2HC1  or  73  g.  It  also  expresses  the 
fact  that  22.4  liters  of  hydrogen  unite  with  22.4  liters  of  chlorine 
to  give  2X22.4  liters  of  hydrogen  chloride,  or  in  general  that  one 
volume  of  hydrogen  and  one  volume  of  chlorine  unite  to  give 
two  volumes  of  hydrogen  chloride,  the  volumes  being  those  of 

50 


Chemical  Equations  51 

the  gases  measured  in  all  cases  under  standard  conditions.  In 
reactions  involving  gases  the  volume  of  each  gas  taken  or 
formed  is  always  shown  by  the  coefficient  in  front  of  its  formula 
in  the  equation  for  the  reaction. 

78.  The  Equation  for  the  Burning  of  Carbon. — Some  free 
elements  like  carbon  are  not  sufficiently  volatile  to  enable  us  to 
find  the  formula  of  the  element  from  measurements  of  the  vapor 
density  of  the  free  element,  and  in  such  a  case  we  use  the  symbol 
of  the  element  in  writing  equations  involving  its  reactions. 
When  carbon  is  burned  we  find  that  12  g.  of  carbon  require  32  g. 
of  oxygen  occupying  a  volume  of  22 .4  liters,  and  producing  44  g. 
of  carbon  dioxide  occupying  also  a  volume  of  22.4  liters.     These 
facts  may  therefore  be  represented  by  the  equation 

C+O2->CO2. 

Here  the  equation  expresses  directly  the  weights  of  carbon  and 
oxygen  which  unite,  as  well  as  the  weight  of  carbon  dioxide 
formed.  At  the  same  time  it  also  shows  that  22.4  liters  of 
oxygen  when  completely  combined  with  sufficient  carbon  gives 
22.4  liters  of  carbon  dioxide,  but  since  the  carbon  is  not  in  the 
gaseous  state  the  equation  does  not  indicate  anything  regarding 
the  volume  of  the  solid  carbon  which  unites  with  the  volume  of 
oxygen  represented  by  the  formula  O2. 

79.  Solving  Problems. — If  we  remember  that  the  equation 
for  the  burning  of  carbon  in  oxygen  is 

C+02->C02 

we  may  make  use  of  the  facts  represented  by  the  equation  in  the 
solution  of  problems  such  as  the  following:  How  many  liters 
of  oxygen  are  required  for  the  burning  of  5  g.  of  carbon?  To 
solve  this  problem  we  first  write  down  the  equation  which  repre- 
sents the  reaction.  This  shows  that  the  quantity  of  carbon 
represented  by  the  symbol  C,  namely,  12  g.,  requires  for  its  com- 
bustion the  volume  of  oxygen  represented  by  the  formula  O2, 
namely,  22.4  liters.  Therefore  5  g.  of  carbon  would  require 
the  volume  determined  by  the  proportion 


52  Introduction  to  General  Chemistry 

where  x  is  the  number  of  liters  of  oxygen  necessary  for  the  com- 
bustion of  5  g.  of  carbon.  In  an  analogous  manner  we  may 
calculate  what  volume  of  carbon  dioxide  is  produced  by  the 
burning  of  a  known  weight  of  carbon. 

We  may  also  calculate  what  weight  of  oxygen  is  required  or 
carbon  dioxide  produced  in  the  burning  of  5  g.  of  carbon.  If 
12  g.  of  carbon  require  32  g.  of  oxygen,  as  our  equation  indicates, 
then  we  have  only  to  solve  the  following  proportions  in  order  to 
find  the  weight  of  oxygen  required  for  5  g.  of  carbon: 


where  y  is  the  required  answer. 

80.  The  Burning  of  Magnesium.  —  Suppose  we  desire  to  find 
by  experiment  the  formula  of  the  product  formed  by  burning 
magnesium  in  oxygen.  It  will  be  recalled  that  the  metal 
magnesium  in  the  form  of  wire  or  powder  burns  with  great  ease 
in  oxygen,  forming  a  white  solid  substance  which  we  have  called 
magnesium  oxide  (28).  We  find  by  experiment  that  10  g.  of 
magnesium  when  burned  yields  i6.6g.  of  magnesium  oxide. 
Let  us  suppose  that  we  have  discovered  by  careful  experiment 
that  magnesium  oxide  contains  only  the  elements  magnesium 
and  oxygen.  The  difference  between  the  weight  of  the  mag- 
nesium oxide  formed  and  the  magnesium  taken  must  represent 
the  weight  of  oxygen  which  has  combined  with  the  log.  of 
magnesium.  This  we  find  to  be  6  .  6  g. 

Suppose  we  know  that  the  symbol  weight  of  magnesium  is 
24.  3  g.  or 


It  is  now  required  to  calculate  the  relative  numbers  of  symbol 
weights  of  magnesium  and  oxygen  that  unite  to  form  magnesium 
oxide.  We  know  that  10  g.  of  magnesium  unite  with  6  .  6  g.  of 
oxygen.  We  may  then  make  the  proportion 

10:  6.6:  124.3:2, 

from  which  we  find  that  2=16.  Therefore  i6g.  of  oxygen 
represented  by  0,  combine  with  the  weight  of  magnesium  repre- 


Chemical  Equations  53 

sented  by  the  symbol  Mg,  and  consequently  we  may  represent 
the  composition  of  magnesium  oxide  by  the  formula  MgO  and 
write  the  equation  for  the  burning  of  magnesium  thus: 

Mg+O-»MgO 

or  better 

2Mg+O2-»2MgO, 

the  latter  equation  having  the  advantage  in  that  it  shows  the 
volume  of  oxygen,  22.4  liters,  as  well  as  its  weight  required  for 
the  burning  of  the  .weight  of  magnesium  represented  by  2Mg. 
But  since  both  magnesium  and  magnesium  oxide  are  solid 
involatile  substances  the  equation  does  not  show  the  volumes  of 
these  solids  entering  into  the  reaction,  as  it  would  in  the  case  of 
gaseous  substances. 

81.  The  Burning  of  Iron.  —  It  will  be  recalled  (17)  that  iron 
burns  in  oxygen,  giving  iron  oxide,  the  formula  for  which  we 
may  now  calculate.  In  an  experiment  in  which  i2.6g.  of  iron 
was  burned  the  weight  of  iron  oxide  produced  was  17.4  g.,  from 
which  we  find,  by  subtracting  the  weight  of  the  iron  burned,  the 
weight  of  the  oxygen  to  be  4.8  g.  These  weights  of  iron  and 
oxygen  must  be  in  the  same  ratio  that  some  number  of  times 
56,  the  symbol  weight  of  iron,  is  to  some  number  of  times  16 
where  these  numbers  are  small  integers.  Dividing  12.6  by  56 
we  get  0.225.  Dividing  4.8  by  1  6  we  get  0.300.  Since  these 
numbers  0.225  and  0.300  are  not  equal,  the  formula  cannot 
be  FeO.  It  will,  however,  readily  be  found  that  0.225  *s  to 
0.300  as  3  is  to  4,  and  therefore  that  12.6:4.8:  13X56:4X16, 
which  shows  that  the  formula  of  the  oxide  of  iron  formed  by 
burning  iron  in  oxygen  is  Fe3O4.  We  may  then  write  the  equa- 
tion for  the  burning  of  iron  as  follows: 


82.  The  Action  of  Hydrogen  on  Copper  Oxide.  —  It  will  be 
remembered  that  we  found  earlier  that  heated  copper  oxide  and 
hydrogen  give  metallic  copper  and  water  (33)  .  In  a  quantitative 
experiment  it  was  found  that  2-387g.  of  copper  oxide  yielded 
1.907  g.  of  copper  and  o.54g.  of  water.  From  the  weights  of 


54  Introduction  to  General  Chemistry 

copper  and  copper  oxide,  together  with  a  knowledge  of  the  fact 
that  copper  oxide  is  composed  of  copper  and  oxygen  only,  we  may 
discover  very  readily  that  the  formula  of  copper  oxide  is  CuO, 
knowing  the  symbol  weight  of  copper  to  be  63 . 6.  Furthermore, 
since  water  contains  only  hydrogen  and  oxygen  and  o .  54  g.  of 
water  has  been  formed  from  2.387—  i  .907  or  0.48  g.  of  oxygen, 
the  weight  of  hydrogen  present  in  the  o.  54  g.  of  water  must  have 
been  o .  06  g.  Making  a  calculation  analogous  to  that  made  in 
finding  the  formula  for  iron  oxide,  we  find  that  o :  06 :  o .  48 : :  2  X 
1:1X16  and  that  therefore  the  composition  of  water  is  repre- 
sented by  the  formula  H2O.  We  may  now  write,  as  the  equation 
for  the  reaction  between  copper  oxide  and  hydrogen, 

CuO+H2->Cu+H20. 

83.  The  Action  of  Acetylene  on  Copper  Oxide. — From  what 
has  preceded  the  student  will  understand  that  in  order  to  be  able 
to  write  the  equation  for  any  reaction  we  must  know  all  of  the 
substances  entering  into  the  reaction  and  all  of  the  products.  In 
addition  we  must  know  the  formula  of  each  substance.  We  may 
illustrate  the  method  then  employed  by  means  of  reaction 
between  acetylene  and  copper  oxide  which  we  have  already 
studied. 

When  acetylene  is  passed  over  heated  copper  oxide  we  obtain 
carbon  dioxide  and  water,  while  metallic  copper  is  left  behind, 
these  three  substances  being  the  sole  products  of  the  reaction  (50). 
The  formula  of  acetylene  is  C2H2  (62).  The  quantity  of  carbon 
represented  by  C2  would  give  the  quantity  of  carbon  dioxide 
represented  by  2C02;  and  the  quantity  of  hydrogen  represented 
by  H2  would  give  the  quantity  of  water  represented  by  H2O,  so 
that  the  quantities  of  carbon  and  hydrogen  represented  by  one 
formula  weight  of  acetylene  C2H2  would  yield  the  quantities  of 
carbon  dioxide  and  water  represented  by  2C02+H2O.  The 
quantity  of  oxygen  contained  in  the  quantities  of  carbon  dioxide 
and  water  represented  by  2CO2-f-H2O  is  represented  by  50, 
which  quantity  is  contained  in  the  amount  of  copper  oxide 
represented  by  5 CuO.  It  will  thus  appear  that  the  quantity 
of  acetylene  represented  by  C2H3  will  require  the  quantity  of 


Chemical  Equations  55 

copper  oxide  represented  by  sCuO,  and  there  will  be  produced 
the  quantities  of  the  three  products  represented  by 

2C02+H2O+5Cu. 
The  equation  is  therefore 

C2H2+sCuO->  2CO2+H20+5Cu. 

84.  The  Action  of  Ammonia  on  Copper  Oxide.  —  In  an  analo- 
gous manner  we  may  obtain  as  the  equation  for  the  reaction 
which  occurs  when  ammonia  gas  is  passed  over  heated  copper 
oxide,  in  which  case  water,  nitrogen,  and  metallic  copper  are 

formed, 

2NH3+3CuO->3H20+3Cu+N3.  (52) 

85.  The  Meaning  of  an  Equation.  —  Since  chemists  make 
extensive  use  of  equations,  it  is  of  fundamental  importance  that 
the  student  should  understand  exactly  how  equations  are  ob- 
tained and  what  they  mean.    In  every  case  before  the  equation 
for  the  reaction  can  be  written  the  reaction  must  have  been 
thoroughly  investigated  by  experiment  in  the  manner  illustrated 
in  the  preceding  examples.    The  equation  then  shows  at  a  glance 
what  substances  enter  into  and  are  formed  as  a  result  of  the 
reaction.    It  also  shows  the  composition  of  each  of  the  substances 
concerned  and  the  proportions  in  which  they  take  part  in  the 
reaction,  it  being  assumed  in  all  cases  that  we  know  the  weight 
for  which  the  symbol  of  each  element  stands. 

86.  An  Equation  Balances.  —  It  is  one  of  the  most  funda- 
mental facts  in  chemistry  that  in  chemical  change  no  material  is 
destroyed  but  that  the  elements  merely  change  their  forms  of 
combination  with  one  another.    This  important  fact,  which  we 
know  as  the  Law  of  the  Indestructibility  of  Matter,  is  also  repre- 
sented in  every  chemical  equation.    For  we  notice  that  in  each 
equation  we  have  on  each  side  the  same  number  of  symbol 
weights  of  each  element.    Thus  in  the  equation 


we  see  that  there  are  on  each  side  two  symbol  weights  of  carbon, 
two  symbol  weights  of  hydrogen,  five  symbol  weights  of  copper, 


56  Introduction  to  General  Chemistry 

and  five  symbol  weights  of  oxygen.     This  fact  is  usually  expressed 
by  saying  that  the  equation  balances. 

All  of  the  reactions  which  we  have  studied  up  to  this  time 
have  been  thoroughly  investigated  by  chemists  and  for  each  the 
reaction  equation  has  been  discovered.  We  may  now  give,  in 
Table  VI,  a  list  of  such  equations  for  purposes  of  reference.  It 
is  not  to  be  expected,  however,  that  the  student  should  make 
great  effort  to  memorize  all  of  these  equations,  although  such  a 
task  would  not  be  very  difficult,  for,  as  a  little  inspection  will 
show,  there  are  certain  regularities  observable  which  make  this 
a  less  difficult  task  than  might  at  first  sight  seem  to  be  the  case. 

TABLE  VI 

EQUATIONS  OF  OTHER  REACTIONS  STUDIED 
2Hg+O2->2HgO 
2H2+O2->2H2O 
2Na+2H2O->  2NaOH+H2 
NaOH+HCl->NaCl+H2O 
2Na+Cl2H>2NaCl 
Mg+HaO-»MgO+Ha 
3Fe+4H20  -»  Fe3O4+4H2 
CH4+2O2->CO2+2H2O 
C3H8+ioCuO->3CO2+4H2O+ioCu 

87.    Problems 

1.  What  weight  of  mercury  can  be  obtained  by  the  decom- 
position of  10  g.  of  mercuric  oxide? 

2.  What  volume  of  oxygen  at  o°  and  76  cm.  can  be  made 
from  8  g.  of  mercuric  oxide  ? 

3.  What  weight  of  sodium  must  be  acted  on  by  water  to 
yield  500  c.c.  of  hydrogen  at  o°  and  76  cm.? 

4.  What  weight  of  common  salt  can  be  made  from  10  g.  of 
metallic  sodium? 

5.  What  volume  of  hydrogen  at  20°  and  72  cm.  would  be 
formed  by  the  action  of  sufficient  steam  on  6  g.  of  magnesium? 

6.  What  weight  of  copper  oxide  would  be  required  for  the 
oxidation  of  200  c.c.  of  propane  measured  at  25°  and  74  cm.? 
(See  last  equation  of  Table  VI  above.) 

What  weight  of  water  would  be  formed? 


CHAPTER  VII 
ACIDS,  BASES,  AND  SALTS— I 

88.  Caustic  Soda  or  Sodium  Hydroxide. — Let  us  now  con- 
sider the  chemical  changes  which  occurred  in  the  formation  of 
common  salt  from  metallic  sodium,  which  we  have  already 
studied  experimentally.     It  will  be  recalled  that  sodium  reacted 
violently  with  water,  giving  hydrogen  and  sodium  hydroxide,  the 
reaction  being  represented  by  the  equation 

2Na+  2H2O  ->  2NaOH+H2.  (40) 

If  we  repeat  the  experiment  and  evaporate  the  water  we  find 
that  sodium  hydroxide  (also  known  as  caustic  soda)  is  left  as  a 
white  solid  which  is  readily  soluble  in  water.  This  solution  feels 
"soapy"  to  the  fingers  and  if  greatly  diluted  with  water  is  found 
to  have  an  unpleasant  "soapy"  taste.  (It  must  not  be  tasted 
unless  greatly  diluted  with  water,  since  the  concentrated  solu- 
tion acts  powerfully  on  the  mucous  membrane.)  A  piece  of  red 
litmus  paper  is  turned  blue  if  dipped  in  the  solution.  We  know 
many  other  substances  which  have  properties  similar  to  those 
of  sodium  hydroxide.  Such  substances  are  called  bases;  they 
also  have  other  characteristic  properties,  the  most  important  of 
which  we  may  now  consider. 

89.  Bases  Neutralize  Acids. — We  have  learned  (41)   that 
caustic  soda  and  hydrochloric  acid  (which  is  a  solution  of  hydro- 
gen chloride  in  water)  react  to  give  common  salt.    The  equation 
for  this  reaction  is 

NaOH+HCl->  H2O-}-NaCl. 

If  we  add  more  than  sufficient  of  the  acid  and  then  evaporate 
the  solution  to  dryness,  the  excess  of  hydrogen  chloride  will  pass 
off  with  the  water  and  nothing  but  pure  salt,  the  chemical  name 
of  which  is  sodium  chloride,  will  remain.  //  we  test  hydro- 
chloric acid  with  blue  litmus  we  find  that  the  latter  is  turned  red, 
even  by  a  very  dilute  solution.  But  we  find  that  a  solution  oj 

57 


58  Introduction  to  General  Chemistry 

pure  common  salt  in  water  has  no  effect  on  either  blue  or  red  litmus: 
it  is  neutral. 

90.  Properties  of  Acids. — If  we  again  add,  drop  by  drop,  a 
solution  of  hydrogen  chloride  to  one  of  sodium  hydroxide  to 
which  a  few  drops  of  a  solution  of  litmus  have  been  added,  we 
find  that  the  change  of  color  from  blue  to  red  is  produced  suddenly 
and  not  gradually,  a  single  drop  being  sufficient  to  cause  the 
change.     If  we  stop  adding  hydrogen  chloride  at  this  point  we 
find  that  the  solution  consists  only  of  pure  salt  and  water  (with 
but  a  minute  amount  of  litmus).     It  no  longer  has  the  taste  of 
the  sodium  hydroxide,  but  only  that  of  salty  water.    A  diluted 
solution  of  hydrogen  chloride  has  a  rather  agreeable  sour  taste, 
reminding  one  of  vinegar  or  lemon  juice.     Our  experiment  has 
shown  that  both  the  taste  and  the  behavior  toward  litmus  of 
sodium  hydroxide  and  hydrogen  chloride  have  been  changed  in 
their  interaction.     We  say  that  they  have  neutralized  each  other. 
We  know  very  many  substances  which  will  neutralize  sodium 
hydroxide;   all  of  these  have  a  sour  taste  and  color  litmus  red. 
We  call  such  substances  acids,  the  common  name  of  hydrogen 
chloride  solution  being  hydrochloric  acid. 

91.  Another  Base;    Ammonium  Hydroxide. — As  we  have 
already  seen  (51),  ammonia  gas  dissolves  re.adily  in  water,  giving 
a  solution  which  turns  litmus  blue,  and  we  are  not  surprised  to 
find  that  it  neutralizes  hydrochloric  acid.     If  we  evaporate  the 
neutralized  solution  we  obtain  a  white  crystalline  substance,  the 
composition  of  which  is  represented  by  the  formula  NH4C1. 
Since  ammonia  gas  has  the  formula  NH3  and  hydrogen  chloride 
the  formula  HC1,  we  might  be  inclined  to  write  the  equation 

NH3+HC1->NH4C1, 

and,  in  fact,  just  this  reaction  takes  place  if  we  bring  the  two 
gases  together,  a  dense  white  cloud  of  the  solid  product  being 
formed.  However,  if  a  very  concentrated  solution  of  ammonia 
in  water  is  cooled  to  a  very  low  temperature,  we  may  obtain 
crystals  of  a  substance  having  a  composition  represented  by  the 
formula  NH4OH  and  called  ammonium  hydroxide.  This  sub- 
stance is  formed  thus: 

NH,+H20-»NH4OH. 


Acids,  Bases j  and  Salts — 7  59 

We  might  think  to  obtain  it  by  the  evaporation  of  the  water 
solution  of  ammonia;  but  instead  we  get  only  ammonia  gas  and 
water  vapor.  In  fact,  the  crystals  of  ammonium  hydroxide 
obtained  at  a  low  temperature  undergo  a  similar  change  if  they 
are  not  kept  very  cold.  We  say  that  ammonium  hydroxide 
dissociates  readily  into  ammonia  and  water.  Chemists  think  that 
in  a  water  solution  of  ammonia  part  of  the  latter  is  combined 
with  water  to  form  ammonium  hydroxide.  It  is  this  substance 
which  is  thought  to  act  directly  on  red  litmus,  changing  it  to 
blue,  and  to  act  on  hydrochloric  acid  as  follows: 

NH4OH+HC1  ->  NH4C1+H2O. 

We  therefore  call  ammonium  hydroxide  a  base. 

92.  Ammonium  Chloride,  Salts. — The  substance  NH4C1  is 
called  ammonium  chloride.     In  appearance,  taste,  and  other 
properties  to  be  studied  later,  sodium  chloride  and  ammonium 
chloride  closely  resemble  one  another.    They  are  examples  of  an 
important  class  of  chemical  substances  called  salts. 

A  review  of  the  two  neutralizations  just  discussed  will  show 
that  they  have  much  in  common :  in  each  case  a  base  reacts  with 
an  acid  to  form  a  salt  and  water.  Somewhat  later,  other  important 
facts  regarding  neutralization  will  be  discovered.  Before  dis- 
cussing such  matters  we  will  first  become  acquainted  with  a  few 
other  important  acids,  bases,  and  salts. 

93.  Sulfuric  Acid. — One  of  the  most  important,  if  not  the 
most  important,  of  all  acids  is  a  substance  which  is  known  as  oil 
of  vitriol  or   sulfuric  acid.     It  is  manufactured  in  immense 
quantities  and  is  very  cheap,  the  commercial  grade  selling  for 
less  than  one  cent  a  pound.    We  shall  not  now  consider  the 
method  of  its  manufacture  further  than  to  state  that  it  is  made 
from  sulfur.     Its  composition  is  represented  by  the  formula 
H2SO4.     It  is  a  colorless  liquid  of  "oily"  consistency,  but  is 
not  really  an  oil,  as  it  will  mix  with  water  in  all  proportions. 
It  must  be  handled  with  caution,  since  it  can  cause  bad  burns 
if  it  is  spilled  on  the  skin.    (In  case  of  accident,  wash  off  the  acid  in 
much  running  water,  immediately.)     When  sulfuric  acid  is  mixed 
with  water,  the  mixture  gets  boiling  hot,  for  which  reason  the  acid 


60  Introduction  to  General  Chemistry 

should  be  added  very  slowly,  with  stirring,  to  the  water,  if  a 
dilute  solution  is  to  be  made. 

94.  Neutralization  of  Sulfuric  Acid,  Sodium  Sulfate. — We 
find  that  the  dilute  solution  has  a  sour  taste  and  that  it  turns 
litmus  red.  We  may  next  try  whether  it  will  neutralize  a  solu- 
tion of  sodium  hydroxide,  for  which  purpose  we  may  add  to  a 
dilute  solution  of  sulfuric  acid  a  few  drops  of  litmus  solution  and 
then  run  in  sodium  hydroxide  solution  drop  by  drop  until  neutral- 
ity is  reached.  If  the  neutral  solution  is  now  boiled  until  a  solid 
begins  to  appear  and  then  is  left  to  evaporate  at  rocrn  tempera- 
ture, large,  transparent,  glassy-looking  crystals  will  be  formed. 
These  crystals  dissolve  readily  in  water  to  form  a  neutral  solu- 
tion, which  does  not  have  a  sour  taste. 

If  we  allow  the  dry  crystals  to  remain  in  the  open  air  we  find 
that  they  lose  weight  rapidly  and  turn  white  upon  the  surface, 
forming  a  fine  white  powder.  Finally  nothing  is  left  of  the  large, 
clear,  glassy  crystals;  only  the  powder  remains,  the  weight  of 
which  is  much  less  than  that  of  the  original  material.  What  is 
the  cause  of  this  curious  change?  Let  us  put  one  of  the  large 
clear  crystals  into  a  dry  test  tube  and  heat  gently  the  lower  end 
of  the  tube  containing  the  crystal,  while  the  tube  is  held  nearly 
horizontally.  We  soon  see  that  water  has  collected  in  large 
amount  in  the  cold  end  of  the  tube,  while  only  a  white  powder 
is  left  behind.  It  is  now  easy  to  understand  what  occurred  when 
the  large  crystal  was  exposed  in  the  open  air.  It  dissociated  into 
the  white  powder  and  water  which  disappeared  as  vapor.  The 
analysis  of  the  thoroughly  dried  powder  would  show  that  it 
contains  only  sodium,  sulfur,  and  oxygen,  and  in  the  proportions 
represented  by  Na2SO4,  and  since  the  clear  crystals  yielded  only 
Na2SO4  and  water,  their  composition  must  be  represented  by 
Na2S04.#H2O,  where  x  is  a  whole  number  which  must  be  found 
by  means  of  a  quantitative  analysis.  We  call  the  original  sub- 
stance the  hydrate  of  sodium  sulfate,  a  hydrate  of  a  salt  being 
a  compound  of  the  salt  with  water. 

We  may  now  make  the  equation  for  the  formation  of  this 
salt  from  sulfuric  acid.  We  took  H2SO4  and  NaOH  and  got 
^,  from  which  we  see  that  if  two  formula  weights  of  water 


Acids,  Bases,  and  Salts — /  61 

were  formed  from  one  formula  weight  of  H2S04  ana  two  of  NaOH, 
the  whole  of  the  material  taken  would  be  accounted  for  thus: 
H2SO4+  2NaOH  ->  Na2SO4-f-  2H2O. 

This  conclusion  is  rendered  probable  by  the  fact  that  in  the  other 
neutralizations  we  have  studied  water  was  always  one  of  the 
products;  it  may  be  confirmed  by  mixing  with  dry  sodium 
hydroxide  pure  sulfuric  acid,  whereupon  water  and  Na2S04  will 
result.  The  salt  Na2SO4  is  called  sodium  sulfate.  Crystals  of 
anhydrous  sodium  sulfate  are  different  in  form  from  those  of  the 
hydrate. 

95.  Quantitative  Analysis  of  a  Hydrate. — Let  us  now  consider 
the  quantitative  composition  of  the  large,  glassy  crystals  which 
yielded  Na2SO4  and  water.     If  we  weigh  a  crystal  contained  in  a 
porcelain  dish  and  allow  it  to  stand  a  day  or  two  at  room  tempera- 
ture we  find  that  only  the  white  powder  remains.     If  we  now  heat 
the  dish  and  contents  over  a  flame  in  order  thoroughly  to  dry 
the  powder,  and  let  it  cool  and  weigh  it  again,  it  is  obvious  that 
the  loss  of  weight  will  represent  the  weight  of  water  originally 
combined  with  the  weight  of  dry  Na2S04  left  in  the  dish. 

96.  Sodium   Sulfate    Decahydrate :     Na2SO4'ioH2O.— Now 
suppose  that  5 . 796  g.  of  the  hydrate  of  sodium  sulfate  yielded 
2. 556  g.  of  dried  sodium  sulfate,  Na2SO4,  what  is  the  formula  of 
the  hydrate?    In  other  words,  what  is  the  numerical  value  of 
x  in  the  formula  Na2S04-#H20?    The  weight  of  water  driven  off 
was    5.796— 2. 556  =  3.240 g.     We    may    therefore    write    the 
proportion,  2.556  is  to  3.240  as  the  formula  weight  of  sodium 
sulfate  is  to  the  x  times  the  formula  weight  of  water.    Now,  the 
formula  weight  of  sodium  sulfate  is  2X23+32+4X16  =  142 
and  that  of  water  is  2X1+16  =  18.    Therefore  2.556:3.240:: 
142 :  i&x,  from  which  we  find  that  #  =  10,  and  are  thus  led  to  the 
conclusion  that  the  hydrate  of  sodium  sulfate  has  the  formula 
Na2SO4*ioH2O.     If   the   reaction   between   sulfuric   acid   and 
sodium   hydroxide   is   represented   by   the   equation  H2S04+ 
2NaOH->Na2SO4+2H20,    then    the    hydrate    Na2SO4-ioH2O 
must  have  resulted  from  the  union  of  the  sodium  sulfate  with 
part  of  the  water  which  formed  the  solution,  thus : 

Na2SO4+  ioHaO  -»  Na2SO4  •  ioH2O. 


62  Introduction  to  General  Chemistry 

This  substance  is  called  sodium  sulfate  decahydrate  (deca  mean- 
ing ten). 

97.  Hydrates. — Sodium    sulfate    forms    other    compounds 
with  water,  namely  Na^CVyHaO  and  Na^SCVHaO;    but  the 
decahydrate   is   the   common   one.     Many   other   salts   form 
hydrates  and  some  form  a  series  of  hydrates,  as  this  salt  does. 
But  it  must  not  be  supposed  that  all  salts  form  hydrates.     For 
example,  sodium  chloride  and  ammonium  chloride  do  not. 

Solutions  of  the  hydrated  salt  have  exactly  the  same  prop- 
erties as  those  of  solutions  of  the  anhydrous  salt. 

98.  Sodium  Hydrogen  Sulfate:    NaHSO4.~ If  we  exactly 
neutralize  a  definite  quantity  of  sulfuric  acid  with  a  solution  of 
sodium  hydroxide,  noting  the  volume  of  the  latter  used,  and  again 
add  to  a  second  portion  of  sulfuric  acid,  exactly  equal  to  the  first, 
exactly  half  as  much  sodium  hydroxide  solution  as  that  used 
in  the  first  case,  we  find  that  the  first  solution  yields  when 
evaporated  pure  sodium  sulfate,  Na2S04 ;  while  the  second  gives 
crystals  having  a  different  shape  and  appearance,  and  different 
chemical  properties.    Analysis  shows  that  the  composition  of 
these  crystals  is  represented  by  the  formula  NaHSO4.    The 
substance  is  called  sodium  hydrogen  sulfate.    The  equation  for 
the  reaction  in  the  second  case  is 

H2SO4+NaOH->  NaHSO4+H2O. 

99.  The  Law  of  Definite  Composition  Again. — We  may  now 
consider  one  of  the  most  important  and  fundamental  of  all 
chemical  questions,  namely,  whether  the  proportions  of  the 
elementary  constituents  of  a  substance  are  dependent  upon  the 
proportions  which  we  take  of  the  substances  from  which  we  form 
the   product   in   question.     For   example,    we   may   inquire 
whether  we  could  get  a  sulfate  of  sodium  with  a  somewhat 
larger  or  smaller  percentage  of  sodium  if  we  had  used,  in  the 
preceding  experiment,   other  proportions  of  acid   and   base. 
Experiment  will  show,  however,  that  if  we  had  added  a  little 
more  or  less  sodium  hydroxide  we  would  still  have  been  able  to 
obtain  much  NaHS04,  but  that  in  such  cases  there  would  also 
be  some  NazSC^  formed  or  a  little  free  sulfuric  acid  left  after  all 


Acids j  Bases,  and  Salts — 7  63 

the  NaHSO4  had  been  separated  from  the  water.  Facts  like 
these  which  are  met  with  on  every  hand  give  a  special  significance 
to  the  Law  of  Definite  Composition. 

100.  Acid  Properties  of  Sodium  Hydrogen  Sulfate. — We  see 
that  sodium  sulfate,  Na2SO4,  contains  exactly  twice  the  weight  of 
sodium  for  a  given  weight  of  sulfur  and  oxygen  as  does  sodium 
hydrogen    sulfate,    NaHSO4.     Moreover,    we    have    become 
acquainted  with  the  important  fact  that  sulfuric  acid  can  form 
two  sorts  of  sodium  salts.    If  we  dissolve  crystals  of  sodium  hydro- 
gen sulfate  in  water,  we  find  that  the  dilute  solution  has  a  sour 
taste  and  it  turns  litmus  red,  for  which  reasons  we  should  be 
inclined  to  say  that  it  has  acid  properties.    In  accord  with  this 
view,  we  find  that  the  solution  will  readily  neutralize  a  solution 
of  sodium  hydroxide,  giving  sodium  sulfate  and  water,  thus: 

NaHSO4+NaOH  ->  Na2SO4+H2O. 

1 01.  Ammonium  Sulfate  "and  Ammonium  Hydrogen  Sul- 
fate.— If  we  completely  neutralize  sulfuric  acid  with  a  solution 
of  ammonium  hydroxide,  we  obtain  a  salt  called  ammonium  sul- 
fate (NH4)2S04,  thus: 

H3SO4+2NH4OH-»  (NH4)2SO4+2H2O; 

while  with  half  the  proportion  of  ammonium  hydroxide  we  obtain 
ammonium  hydrogen  sulfate,  thus: 

H2SO4+ NH4OH  ->  NH4HSO4+H2O. 

102.  Monobasic  and  Dibasic  Acids:  Acid  Salts  and  Neutral 
Salts. — Hydrochloric  acid  reacts  with  sodium  hydroxide  only  in 
one  proportion,  thus: 

HCl+NaOH  ->  NaCl+H2O, 

for  which  reason  we  call  it  a  monobasic  acid;  but  since  one 
formula  weight  of  sulfuric  can  unite  with  a  maximum  of  two 
formula  weights  of  sodium  hydroxide  we  call  sulfuric  acid  a 
dibasic  acid.  Salts  in  which  but  half  the  maximum  quantity 
of  base  has  been  neutralized  are  usually  called  acid  salts,  because 
they  still  have  acid  properties.  Thus  we  frequently  speak  of 
sodium  acid  sulphate,  meaning  NaHS04.  Chemists  know  many 


Introduction  to  General  Chemistry 


other  dibasic  acids,  all  of  which  also  can  form  acid  salts  as  well  as 
neutral  salts,  as  salts  like  NaaSC^  are  called. 

103.  Making  Hydrochloric  Acid  from  Common  Salt. — If  we 
place  in  a  flask  (Fig.  23)  58  g.  of  dry  common  salt  and  100  g.  of 
sulfuric  acid,  to  which  30  g.  of  water  have  been  added,  and  warm 
the  mixture,  a  change  occurs  with  the  production  of  a  colorless 
gas  which  dissolves  in  water  very  readily,  giving  a  solution  which 
we  can  easily  recognize  as  hydrochloric  acid.    After  the  action 
of  the  sulfuric  acid  on  the  salt  is  complete,  a  white  solid  is  left  in 
the  flask,  which  may  easily  be  dissolved  in  water.     By  evaporat- 
ing part  of  the  water,  and  letting  the  solution  stand  a  while,  we 

/!==__====^  may  obtain  colorless,  transparent 

crystals  of  sodium  hydrogen 
sulfate.  The  following  equation 
represents  the  reaction: 

NaCl+H2SO4  ->  NaHSO4+HCl. 

We  have  to  deal  here  with  a  new 
__  sort  of  chemical  change — one  in 
which  an  acid  acts  upon  a  salt  of 
another  acid  to  give  a  salt  of  the 

first  acid  and  to  produce  the  acid  corresponding  to  the  first  salt. 

This  is  a  very  important  kind  of  chemical  reaction,  which  we 

shall  frequently  make  use  of,  since  by  its  means  we  may  make 

acids  from  their  salts. 

104.  Making  Nitric  Acid  from  Chile  Saltpeter. — We  shall  now 
use  the  method  just  described  for  the  preparation  of  a  new  acid 
from  a  white,  crystalline  substance  called  Chile  saltpeter,  which 
is  found  in  large  quantities  as  a  mineral  substance  in  the  desert 
region  of  Chile. 

If  we  place  85  g.  of  Chile  saltpeter  in  a  retort  (Fig.  24),  add 
100  g.  of  sulfuric  acid,  mixed  with  30  c.c.  of  water,  and  then  heat 
the  mixture  gently,  a  yellow-colored  liquid  may  be  collected  in  a 
cooled  flask.  This  yellow  liquid  gives  off  a  brown  gas  and 
becomes  colorless  when  boiled  a  few  minutes.  Its  analysis  shows 
its  formula  to  be  HN03  and  it  is  called  nitric  acid.  It  is  a  color- 
less liquid  which  may  be  boiled  and  distilled  in  glass  vessels. 


FIG.  23 


Acids,  Bases,  and  Salts — /  65 

Pure  or  concentrated  nitric  acid  is  even  more  dangerous  than 
sulfuric  acid,  causing  serious  burns  and  destroying  clothing, 
and  must  be  handled  with  greatest  care.  It  will  mix  with  water 
in  all  proportions,  giving  a  solution  which,  when  very  dilute,  has 
a  sour  taste  and  turns  litmus  red. 

When  nitric  acid  is  mixed  with  sodium  hydroxide  solution  the 
latter  is  neutralized,  a  salt  of  the  composition  NaNO3  and  water 
being  the  only  products,  as  represented  by  the  equation 

HN03+NaOH  -»  NaNO3+H2O. 

The  salt,  which  is  called  sodium  nitrate,  is  found  to  be  identical 
with  purified  Chile  saltpeter.  The 
action  of  sulfuric  acid  on  saltpeter 
leaves  in  the  retort  a  white  solid 
which  closely  resembles  that  left 
when  salt  is  heated  with  sulfuric  acid, 
and,  in  fact,  the  residue  is  easily  found 
to  be  the  same  substance,  sodium 
hydrogen  sulfate,  NaHS04.  The 
equation  for  the  reaction  is  therefore  24 

NaNO3+H2SO4->  NaHSO4+HNO3. 

105.  The  Action  of  Nitric  Acid  on  Ammonium  Hydroxide. — 

We  may  now  propose  a  question  to  be  answered,  not  after  direct 
experiment,  but  as  a  result  of  the  general  knowledge  we  have 
gained  regarding  the  behavior  of  the  acids  and  bases  already 
studied.  It  is:  What  would  be  the  result  of  mixing  nitric  acid 
and  ammonium  hydroxide?  We  recall  that  hydrochloric  acid 
and  sodium  hydroxide,  a  base,  give  sodium  chloride  and  water, 

thus: 

HCl+NaOH->  NaCl+H20; 

that  the  same  acid  gives  with  ammonium  hydroxide,  also  a  base, 
ammonium  chloride  and  water,  thus: 

HC1+NH4OH  ->  NH4C1+H2O. 

Furthermore,  we  have  just  seen  (104)  that  nitric  acid  and  sodium 
hydroxide  give  sodium  nitrate  and  water,  thus : 
HNO3+NaOH  ->  NaNO3+H2O, 


66  Introduction  to  General  Chemistry 

and  we  would  certainly  expect  that  nitric  acid  and  ammonium 
hydroxide  would  behave  analogously  and  give  ammonium  nitrate 
and  water,  thus: 

HNO3-f-NH4OH  ->  NH4NO3+H2O. 

Now  this  is  precisely  what  takes  place  when  we  test  our  prediction 
by  experiment.  We  seem,  therefore,  to  have  discovered  the 
secret  of  the  way  in  which  acids  and  bases  act  toward  each  other. 
It  may  be  summed  up  in  the  statement,  An  acid  and  a  base 
neutralize  each  other,  forming  a  salt  and  water. 

106.  A  New  Base:  Caustic  Potash  or  Potassium  Hydroxide. 
—Let  us  now  take  up  the  study  of  a  new  base,  caustic  potash, 
which  closely  resembles  caustic  soda  (sodium  hydroxide).  It 
will  be  remembered  that  the  metal  sodium  reacts  violently  with 
water,  giving  sodium  hydroxide  and  hydrogen  gas,  thus: 

2Na+  2HaO  ->  2NaOH+H2.  (88) 

Now,  chemists  know  another  metallic  element,  potassium,  which 
closely  resembles  sodium.  Like  sodium,  it  is  a  silver-white 
metal,  soft  enough  to  be  cut  easily  with  a  knife  and  tarnishing 
very  rapidly  in  the  air.  For  a  reason  that  we  shall  soon  learn  it 
is  kept  covered  with  oil  in  a  carefully  stoppered  bottle.  If  we 
throw  a  small  bit  of  potassium  into  a  beaker  of  water,  it  bursts 
into  a  flame  of  lavender  color,  spinning  and  darting  to  and  fro 
on  the  surface  of  the  water  and  completely  disappearing  in  a  few 
moments.  Examination  of  the  water  shows  that  it  will  turn 
litmus  blue,  that  it  has  a  "soapy"  taste,  like  a  very  dilute  solu- 
tion of  sodium  hydroxide,  and  that  a  white  solid  is  left  when  the 
solution  is  evaporated  to  dryness.  This  solid  is  found  by  suitable 
methods  of  analysis  to  contain  the  elements  potassium,  oxygen, 
and  hydrogen  in  the  proportion  represented  by  the  formula 
KOH,  and  is  called  potassium  hydroxide. 

If  instead  of  throwing  the  bit  of  potassium  on  the  surface  of 
the  water  we  bring  it  under  the  mouth  of  an  inverted  cylinder 
filled  with  water,  with  the  mouth  immersed  in  a  vessel  of  water, 
the  potassium  rises  to  the  top  of  the  water  in  the  cylinder,  pro- 
ducing a  gas  which  displaces  the  water  in  the  cylinder,  but  does 


Acids,  Bases,  and  Salts — /  67 

not  take  fire.  The  gas  is  easily  identified  as  hydrogen,  while  the 
water  contains  dissolved  potassium  hydroxide  as  before.  The 
equation  for  the  reaction  in  the  cylinder  is 

•      2K+2H2O->2KOH+H2. 

When  the  action  takes  place  in  the  open  beaker,  the  heat  pro- 
duced sets  fire  to  the  hydrogen,  which  burns,  together  with  a  small 
portion  of  the  potassium. 

107.  Potassium  Salts. — On  account  of  the  behavior  of  a  solu- 
tion of  potassium  hydroxide  toward  litmus  and  also  because  of  its 
"soapy"  feel  and  taste,  we  should  conclude  that  it  is  a  base  and 
if  so  that  it  should  form  salts  with  acids.  We  might  even  venture 
to  predict  the  formulae  of  the  salts  it  would  be  expected  to  form 
with  hydrochloric,  sulfuric,  and  nitric  acids,  and  to  write  the 
equations  as  follows: 

HCl+KOH->  KCl-f-H2O 
H2SO4+  2KOH  ->  K2SO4+  2H2O 
H2SO4+KOH-»  KHSO4+H2O 
HNO3+KOH  ->  KNO3+H2O. 

And  in  every  case  these  predictions  would  be  found  by  experiment 
to  be  correct!  The  potassium  salts  so  formed  are  all  white 
crystalline  solids  and  are  all  soluble  in  water.  All  except  potas- 
sium hydrogen  sulfate  give  solutions  which  are  neutral  to  litmus, 
while  this  salt  has  acid  properties  like  those  of  sodium  hydrogen 
sulfate. 


CHAPTER  VIII 
WATER  AND  SOLUTIONS 

1 08.  Water. — We  have  already  learned  that  pure  water  is 
readily  obtained  by  the  distillation  of  natural  waters  (23),  and 
that  it  is  a  compound  of  hydrogen  and  oxygen,  the  composition 
of  which  is  represented  by  the  formula  H2O  (70).    In  describing 
a  substance  we  shall  often  mention  its  physical  and  chemical 
properties.    The  properties  of  a  substance  embrace:   the  state 
(whether  solid,  liquid,  or  gaseous);  crystalline  form,  if  solid; 
specific  gravity  or  density;  color;  odor;  taste;  conductivity  for 
heat  and  electricity;   boiling-point;    freezing-point,  etc.    The 
chemical  properties  of  a  substance  are  those  which  it  exhibits  in 
its  typical  chemical  reactions. 

109.  The  Physical  Properties  of  Water:   Color.— We  know 
that  according  to  the  temperature  water  can  exist  as  solid,  liquid, 
or  gas.    The  color  of  liquid  water  is  a  very  faint  blue;  so  faint,  in 
fact,  that  it  cannot  be  noticed  in  a  glass  of  water,  but  is  obvious 
in  a  white  bathtub  full  of  clear  water.    The  color  of  large  bodies 
of  clear  water  is  usually  blue,  but  it  may  be  of  any  other  shade 
if  dissolved  or  suspended  impurities  (mud)  are  present.    The 
yellow  color  of  the  waters  of  many  rivers  is  due  to  suspended 
clay;  such  water  is  not  clear,  but  muddy  or  turbid.     Streams  and 
lakes  in  hemlock  forests  often  contain  perfectly  clear  water 
having  the  color  of  tea,  due  to  coloring-matter  dissolved  from 
the  hemlock.    The  clear  green  color  of  some  waters  is  usually 
the  result  of  the  blending  of  the  natural  blue  color  of  the  water 
with  the  yellow  light  reflected  from  the  sand  beneath. 

no.  Specific  Gravity  or  Density. — At  the  temperature  of 
4°  C.,  i  c.c.  of  water  weighs  i  g.  Since  the  specific  gravity  or 
density  of  any  substance  may  be  defined  as  the  weight  of  i  c.c.,  it 
follows  that  water  has  a  specific  gravity  of  i .  ooo  at  4°  C.  Or,  we 
may  say  that  the  specific  gravity  or  density  of  a  substance  is 
found  by  dividing  its  weight  by  the  weight  of  an  equal  volume  of 
water.  Water  has  its  greatest  density  at  4°;  if  a  given  volume 

68 


Water  and  Solutions  69 

of  water  at  4°  is  either  heated  or  cooled,  it  expands  and  therefore 
decreases  in  density. 

in.  Specific  Heat. — The  quantity  of  heat  required  to  raise 
the  temperature  of  i  g.  of  water  i°  C.  is  by  definition  called  one 
calorie.  Water  is  said  to  have  a  specific  heat  of  one  or  unity. 
The  specific  heat  of  any  substance  is  the  quantity  of  heat  in  calories 
required  to  raise  the  temperature  of  one  gram  of  it  one  degree. 
Nearly  all  substances  have  specific  heats  less  than  unity. 

112.  Vapor  Pressure. — Water  contained  in  an  open  vessel 
evaporates  at  all  temperatures,  but  the  more  rapidly  in  propor- 
tion as  the  temperature  is  higher,  other  things  being  equal.     If 
water  evaporates  into  an  evacuated  space  the 
pressure  within  the  space  increases  to  a  value 
which  is  dependent  only  upon  the  tempera- 
ture,   being   greater   in   proportion   as    the 
temperature  is  higher.    The  pressure  so  pro- 
duced is  called  the  vapor  pressure  of  water; 
it  may  easily  be  demonstrated  by  means  of  a 
barometer  tube  filled  with  mercury.    If  we 
prepare  two  such  tubes  (Fig.  25)  and  intro- 
duce a  few  drops  of  water  into  one  by  means 
of  a  suitably  shaped  glass  tube,  the  water  FIG.  2S 

will  rise  until  it  floats  on  the  surface  of  the 
mercury.  At  the  same  time  the  level  of  the  mercury  will 
fall  2  or  3  cm.,  showing  that  a  pressure  has  been  produced 
above  the  mercury  in  the  space  which  has  been  a  vacuum. 
If  the  tube  into  which  the  water  is  introduced  has  a  glass  jacket 
into  which  warm  water  can  be  poured,  it  will  be  found  that 
the  higher  the  temperature  is,  the  higher  the  vapor  pressure  will  be. 
If  we  should  raise  the  temperature  to  100°,  the  level  of  the 
mercury  in  the  barometer  tube  would  sink  to  that  of  the  surface 
of  the  mercury  in  the  dish  in  which  the  tube  stands,  thus  show- 
ing that  the  vapor  pressure  at  100°  is  equal  to  the  pressure  of  the 
atmosphere.  Table  VII  shows  the  vapor  pressure  of  water  at 
various  temperatures  between  o°  and  100°. 

When  the  atmospheric  pressure  is  760  mm.,  water  boils  at 
100°.  Now,  we  see  from  the  table  that  at  100°  the  vapor 


Introduction  to  General  Chemistry 


pressure  is  760  mm.,  therefore  the  boiling-point  is  that  temperature 
at  which  the  vapor  pressure  becomes  just  equal  to  the  normal  atmos- 
pheric pressure,  760  mm.,  which  is  the  average  pressure  at  sea- 


TABLE  VII 


Temperature 

Pressure 

Temperature 

Pressure 

0° 

4  6  mm 

60° 

140  2  mm 

IO 

92 

7O 

272     g 

2O 

17   4. 

80 

see    t 

3O 

31  6 

QO 

$26  o 

40  

CCO 

00 

7  •2-7    2 

•»o.  . 

O2    2 

IOO 

760  o 

level.  At  higher  altitudes,  at  which  the  atmospheric  pressure 
is  less  than  760  mm.,  water  boils  at  temperatures  lower  than 
100°.  Thus  if  the  pressure  is  733 . 2  mm.,  the  boiling-point  is 
99°.  Since  the  atmospheric  pressure  at  a  given  place  is  variable 
through  a  range  of  20  mm.  or  more,  the  boiling-point  at  this 
place  is  not  constant,  but  varies  with  the  rise  and  fall  of  the 
barometer. 

113.  Correction  of  the  Volume  of  a  Gas  for  Vapor  Pressure. — 
Gases  like  hydrogen  and  oxygen,  which  are  not  very  soluble  in 
water,  are  often  measured  in  tubes  in  which  the 
gases  are  confined  by  means  of  water.  Such  gases 
always  contain  water  vapor,  and  part  of  the  total 
pressure  exerted  by  the  gas  is  due  to  the  vapor 
pressure  of  the  water.  The  part  of  the  pressure 
(partial  pressure)  exerted  by  the  gas  itself  is  found 
by  subtracting  from  the  total  pressure  the  vapor 
pressure  of  the  water.  For  example,  suppose  that 
some  hydrogen  is  collected  over  water  in  a  grad- 
uated glass  tube  (Fig.  26).  If  the  position  of  the 
tube  is  adjusted  so  that  the  level  of  the  water  is  the  same  inside 
the  tube  as  outside,  the  total  pressure  within  must  be  exactly 
equal  to  the  atmospheric  pressure,  as  shown  by  the  barometer. 
Suppose  that  the  barometric  pressure  is  748.6mm.  and  the 
temperature  20°.  Table  VII  shows  that  at  20°  the  vapor 
pressure  of  water  is  17.4  mm.,  therefore  the  pressure  due  to  the 


FIG.  26 


Water  and  Solutions  71 

hydrogen  is  748 .6— 17. 4  =  73 1.2  mm.  If  the  observed  volume 
was  30  c.c.,  the  volume,  F,  at  standard  conditions  would  be 

^30X731-2X273^ 

760X293 

114.  Vapor  Pressure  of  Liquids  and  Solids  in  General. — 

Liquids  in  general  readily  pass  into  the  form  of  vapor,  and  just 
as  in  the  case  of  water,  a  given  pure  liquid  has,  at  each  tempera- 
ture, a  definite  vapor  pressure;  but  the  vapor  pressure  of  one 
liquid — say  alcohol — is  not  in  general  the  same  at  a  given 
temperature  as  that  of  another  liquid — say  water.  In  every 
case,  however,  the  boiling-point  of  the  liquid  is  that  temperature  at 
which  its  vapor  pressure  equals  760  mm.  Many  solids,  for 
example,  camphor  and  naphthalene  (moth-balls),  have  appre- 
ciable vapor  pressures  at  room  temperature;  but  the  vapor 
pressures  of  most  solids  at  such  temperature  are  too  small  to  be 
noticeable. 

115.  Latent  Heat  of  Evaporation. — If  it  is  true  that  water 
boils  at  1 00°  because  at  this  temperature  the  vapor  pressure  of 
water  just  equals  the  normal  atmospheric  pressure,  it  may  be 
asked  why  the  whole  of  the  water  does  not  change  at  once  into 
steam  as  soon  as  its  temperature  is  raised  to  100°.    We  know,  of 
course,  that  this  does  not  occur,  and,  further,  that  the  rapidity 
with  which  water  boils  away  is  greater,  the  greater  the  amount 
of  heat  applied.    The  explanation  is  found  in  the  fact  that  it 
requires  a  large  amount  of  heat  to  change  water  at  100°  into  steam 
at  the  same  temperature.    In  fact,  540  calories  of  heat  are  required 
for  the  conversion  of  i  g.  of  water  at  100°  into  steam.    The  heat 
so  used  up  does  not  raise  the  temperature  of  the  substance.     It  is 
consumed  in  changing  the  liquid  water  into  the  gaseous  state; 
it  is  said  to  become  latent,  and  in  consequence  we  say  that  the 
latent  heat  of  evaporation  of  water  is  540  calories.    Every  pure 
liquid  has  a  latent  heat  of  evaporation.    This  differs  from  one 
substance  to  another. 

1 1 6.  Use  of  Steam  for  Heating. — When  steam  cools  to  100°  it 
begins  to  condense  to  liquid  water,  and  for  every  gram  of  steam 
that  condenses  540  calories  of  heat  are  given  out.    The  heat 


72  Introduction  to  General  Chemistry 

so  given  out  may  be  considered  to  be  that  which  became  latent 
when  the  water  was,  by  being  heated,  converted  into  steam.  // 
is  on  account  of  the  latent  heat  given  out  upon  condensation  that 
steam  is  so  effective  in  the  heating  of  buildings:  every  gram  of 
steam  that  condenses  in  the  radiator  liberates  540  calories  of 
heat.  Of  course,  the  further  cooling  of  the  water  in  the  radiator 
gives  out  some  additional  heat. 

117.  Burns  Produced  by  Steam. — It  is  a  well-known  fact 
that  serious  burns  result  when  steam  comes  in  contact  with  the 
skin.    At  first  thought,  this  result  seems  to  be  out  of  harmony 
with  the  fact  that  air  at  100°  can  be  borne  by  the  hand  without 
discomfort.     The  explanation  of  this  difference  is  found  in  the 
fact  that  gases  (including  the  vapors  of  boiling  liquids)  are  very 
poor  conductors  of  heat  as  compared  with  liquids.     Steam  at  100° 
partly  condenses  on  striking  the  skin  and  wets  it  with  a  layer 
of  boiling-hot  water,  which  is  a  good  conductor  of  heat.     Further- 
more, since  540  calories  of  heat  are  given  out  by  every  gram  of 
steam  condensed  to  water,  the  latter  is  kept  at  100°  as  long  as 
steam  is  present.     On  the  other  hand,  air  is  so  poor  a  conductor 
of  heat  that  the  skin  is  not  burned  by  a  brief  exposure  to  it  at  100°. 

118.  Latent  Heat  of  Fusion  of  Ice. — Ice  melts  at  o°;  but  all 
of  a  given  mass  of  ice  does  not  melt  immediately  when  its 
temperature  is  raised  to  zero.     Just  as  heat  is  required  to  change 
liquid  water  into  vapor,  so  also  heat  is  needed  to  change  ice  at 
zero  into  water  at  the  same  temperature.     The  heat  so  absorbed 
is  called  the  latent  heat  of  fusion  of  ice.     It  requires  79  calories 
to  melt  i  g.  of  ice;   therefore  the  latent  heat  of  fusion  of  ice  is 
79  calories.     Every  solid  has  a  definite  and  characteristic  latent 
heat  of  fusion. 

119.  The  Density  of  Ice. — The  density  or  specific  gravity  of 
ice  is  o .  91 7.     It  is  for  this  reason  that  ice  floats  on  water.    The 
expansion  which  occurs  when  water  freezes  exerts  very  great 
pressure,  illustrations  of  which  are  often  seen  in  the  bursting  of 
water  pipes  and  other  vessels  when  water  freezes  in  them.     Not 
all  liquids  expand  upon  freezing;    in  many  cases  contraction 
occurs,  thereby  giving  rise  to  solids  which  sink  in  the  correspond- 
ing liquids. 


Water  and  Solutions  73 

120.  Solutions  and  Suspensions. — The  mixture  which  results 
upon  dissolving  salt  in  water  is  called  a  solution  of  salt  in  water. 
The  terms  "dissolve"  and  "  solution"  are  used  in  chemistry  with 
definite  meanings.     If,  upon  mixing  a  solid  with  a  liquid,  the 
former  partly  or  wholly  disappears  and  the  resulting  liquid  is 
still  clear  and  transparent  and  not  cloudy  or  muddy,  and  if, 
moreover,  upon  allowing  the  liquid  to  evaporate  we  regain  the 
unchanged  solid  substance,  we  say  that  the  solid  had  dissolved 
in  the  liquid  to  form  a  solution.     Either  or  both  of  the  substances 
may  be  colored  and  still  a  clear  (although  colored)  solution  may 
result.     The  liquid  in  which  a  substance  is  dissolved  is  called  the 
solvent. 

If  we  stir  up  some  common  clay  with  water,  much  of  the  clay 
fails  to  settle  out  of  the  water  at  once,  and  we  get  a  cloudy  or 
muddy  fluid,  like  the  water  of  a  muddy  river.  In  this  case  we 
do  not  say  that  the  clay  has  dissolved  in  the  water  or  that  we  have 
a  true  solution  of  the  clay.  We  say  that  the  clay  is  suspended 
in  the  water,  and  call  the  muddy  water  a  suspension.  Clay 
suspended  in  water  will  settle  out  very  slowly  and  finally  leave 
clear  water  above  a  layer  of  mud. 

121.  The  Concentration  of  Solutions. — A  solution  containing 
a  small  proportion  of  a  dissolved  substance  is  said  to  be  dilute, 
while  one  containing  a  large  proportion  is  called  concentrated. 
We  dilute  a  concentrated  solution  by  adding  solvent  to  it,  and 
concentrate  a  dilute  solution  by  evaporating  the  solvent.     We  use 
the  term  concentration  in  discussing  the  relative  amount  of 
dissolved  substances  in  a  solution. 

122.  Solubility  of  Substances:    Saturated  Solutions.— It  is 
easy  to  discover  that  the  amount  of  a  substance  which  will  dis- 
solve in  a  given  amount  of  water,  say  100  c.c.,  depends  upon  the 
nature  of  the  substance  and  upon  the. temperature.     If  we  mix 
some  common  salt  with  about  double  its  weight  of  water  and 
stir  or  shake  the  mixture  a  sufficient  length  of'  time  (usually  one 
to  two  hours),  keeping  the  temperature  constant  all  the  while, 
and  then,  after  allowing  any  suspended  crystals  to  settle,  draw 
off  a  portion  of  the  clear  solution,1  weigh  it,  and  evaporate  the 
water,  we  get  the  salt  dissolved  in  the  portion  of  the  solution 


74  Introduction  to  General  Chemistry 

taken.  By  weighing  the  salt  we  can  readily  find  the  weight  of 
salt  dissolved  in  a  given  weight  of  water  at  the  temperature  at 
which  the  experiment  was  made.  We  find  in  this  way  that 
100  g.  of  water  at  25°  dissolves  37 . 6  g.  of  salt. 

To  make  such  a  solubility  determination  we  must  observe 
several  precautions:  First,  the  amount  of  solid  substance  must 
be  considerably  greater  than  the  amount  of  water  taken  will  dis- 
solve ;  secondly,  the  shaking  must  be  continued  as  long  as  more 
substance  dissolves — this  is  easily  ascertained  by  prolonging  the 
shaking  and  making  additional  determinations  of  the  concentra- 
tion of  the  solution;  thirdly,  the  temperature  must  be  kept 
constant. 

A  solution  which  at  a  fixed  temperature  will  dissolve  no  more 
of  a  given  substance  is  called  a  saturated  solution.  When  we 
speak  of  the  solubility  of  a  substance  we  mean  the  amount  of 
substance  dissolved  in  a  given  amount  of  water  in  the  case  of 
the  saturated  solution.  The  following  brief  table  gives  the 
solubilities  in  water  at  25°  of  several  salts. 

TABLE  VIII 

GRAMS  or  SUBSTANCE  IN  100  G.  OF  WATER  AT  25° 


NaCl.  . 


Na2SO4  •  ioH2O       27 
NaNO3 92 


37  g- 


KC1 |       34g. 

K2SO4 '.        12 

KN03 37 


123.  Supersaturated  Solutions. — At  25°  100  g.  of  water  will 
dissolve  27  g.  of  sodium  sulfate  decahydrate,  Na2SO4-ioH20, 
while  at  30°  the  same  amount  of  water  will  dissolve  40  g.  of  the 
salt.  If  we  make  a  saturated  solution  of  the  salt  at  30°,  having 
an  excess  of  crystals  of  the  salt  present,  and  then  cool  the  whole 
to  25°,  and  keep  it  at  25°,  stirring  or  shaking  it  for  an  hour  or  two, 
more  solid  is  deposited  and  there  results  a  solution  which  contains 
just  the  same  weight  of  the  salt  in  100  g.  of  water  as  a  saturated  solu- 
tion at  25°,  namely,  27  g. 

A  slight  change  in  the  procedure  gives  a  very  different  result 
and  brings  to  light  a  new  phenomenon.  If  the  solution  of  sodium 
sulfate  which  is  saturated  at  30°  is  freed  from  every  particle  of  the 


Water  and  Solutions  75 

solid  crystalline  substance  and  then  allowed  to  cool  to  25°  or  even 
lower,  without  being  stirred  or  shaken,  it  remains  perfectly  clear  and 
does  not  deposit  any  crystals.  Such  a  solution  contains  at  25° 
much  more  sodium  sulfate  than  a  saturated  solution  prepared 
at  25°  in  the  manner  described  in  the  preceding  paragraph. 
This  more  concentrated  solution  is  called  a  supersaturated 
solution.  If  we  now  drop  into  the  supersaturated  solution  a 
crystal  of  sodium  sulfate  (and  for  this  purpose  an  almost  in- 
visible fragment  of  the  crystalline  dust  will  be  sufficient),  the 
formation  of  crystals  will  begin  at  once  and  proceed  until  the 
amount  of  dissolved  substance  per  100  g.  of  water  is  reduced 
exactly  to  that  of  a  saturated  solution  at  the  existing  tempera- 
ture. 

Experience  has  shown  that  a  supersaturated  solution  can 
only  be  obtained  in  the  complete  absence  of  the  solid  substance, 
and  that  a  supersaturated  solution  begins  to  deposit  its  excess 
of  dissolved  substance  when  a  crystal  of  this  same  substance  is 
brought  into  the  solution.  The  deposition  of  crystals  by  a 
supersaturated  solution  can  also  often  be  started  by  shaking 
or  stirring  the  solution  or  by  adding  a  crystal  of  another  sub- 
stance having  the  same  crystalline  form. 

Not  all  substances  form  supersaturated  solutions  equally 
readily.  The  presence  of  impurities  favors  supersaturation. 
Syrups,  preserves,  and  honey  are  often  supersaturated  with 
respect  to  the  sugar  dissolved  in  the  water  present.  When  such 
solutions  "turn  to  sugar,"  this  is  only  the  crystallization  of  the 
excess  of  sugar  above  that  required  to  make  a  saturated  solution. 

124.  Solubility  of  Liquids  in  Liquids. — It  is  proverbial  that 
"oil  and  water  will  not  mix."  On  the  other  hand,  some  pairs  of 
liquids  will  mix  completely  in  all  proportions;  examples  of  such 
combinations  are  water  and  alcohol  and  water  and  sulfuric  acid. 
We  know  other  pairs  of  liquids  that  will  not  dissolve  one  another 
in  all  proportions,  but  that  will  dissolve  one  another  partially. 
Water  and  ether  belong  to  this  class;  100  c.c.  of  water  will 
dissolve  8  c.c.  of  ether,  and  100  c.c.  of  ether  will  dissolve  3  c.c. 
of  water.  If  we  pour  ether  into  water,  we  find  that  the  former 
floats  on  the  surface  of  the  latter.  If  equal  volumes  of  ether  and 


76  Introduction  to  General  Chemistry 

water  are  thoroughly  shaken  together,  the  former  soon  separates 
from  the  latter,  and  two  distinct  layers  result  as  before.  If,  now, 
we  examine  each  layer,  we  find  that  the  water  contains  some 
dissolved  ether  and  the  ether  some  dissolved  water.  This  is  a 
case  of  partial  miscibility. 

125.  Solubility   of    Gases   in   Liquids. — We   have   already 
learned  that  hydrogen  chloride  (44)  and  ammonia  (51)  are  both 
very  soluble  in  water.    At  o°  water  dissolves  550  times  its  own 
volume  of  the  first  gas  and  1,150  times  its  volume  of  the  second. 
No  gas  which  we  have  studied  is  completely  insoluble  in  water; 
for  example,  100  c.c.  of  water  dissolves  2.  i  c.c.  of  hydrogen  and 
4.8  c.c.  of  oxygen.     Fishes  depend  for  their  existence  upon  the 
oxygen  dissolved  in  water;    by  means  of  their  gills  they  take 
from  the  water  the  oxygen  they  require. 

126.  Henry's  Law. — The  solubility  of  all  gases  decreases  with 
rise  of  temperature.    At  a  fixed  temperature  the  weight  of  gas 
dissolved  by  a  given  volume  of  water  of  other  liquid  is  dependent 
upon  the  pressure  of  the  gas  and  is,  in  general,  directly  propor- 
tional to  the  pressure.    This  statement  is  known  as  Henry's 
Law.    The  law  does  not  apply  to  very  soluble  gases,  like  am- 
monia, dissolving  in  water — probably  because  chemical  union 
occurs,  since  we  know  that  NH4OH  is  formed  in  this  case  (91). 

127.  Heat  of  Solution. — If  we  shake  some  potassium  nitrate 
or  ammonium  chloride,  or  indeed  any  one  of  many  salts,  with 
water,  we  find  that  as  the  substance  dissolves  the  solution  becomes 
appreciably  colder.    This  indicates  that  heat  is  required  to  change 
the  solid  into  the  dissolved  state.    This  phenomenon  is  analogous 
to  that  met  with  when  a  solid,  like  ice,  melts.     It  requires  79 
calories  to  melt  i  g.  of  ice,  while  115  calories  are  absorbed  when 
i  g.  of  potassium  nitrate  dissolves.    That  is,  we  must  supply 
115  calories  to  i  g.  of  the  salt,  and  sufficient  water,  in  order  to 
prevent  a  fall  of  temperature  when  solution  takes  place.    The 
heat  so  required  is  called  heat  of  solution. 

When  any  substance  whatever  melts,  heat  is  required,  or  is 
absorbed,  and  we  might  expect,  similarly,  that  heat  will  always 
be  absorbed  when  a  substance  dissolves;  but  this  is  not  the  case. 
Many  substances,  upon  dissolving,  give  out  heat.  In  the  case  of  a 


Water  and  Solutions  77 

few  substances  the  absorption  or  evolution  of  heat  upon  dissolv- 
ing is  very  small.  Common  salt  dissolves  in  water  with  very 
small  heat  absorption. 

128.  Boiling-Point  of  Solutions. — It  is  very  easy  to  show  that 
a  solution  of  a  solid  substance,  like  salt  or  sugar  in  water,  boils 
at  a  higher  temperature  than  pure  water.     This  is  an  invariable 
rule  for  solutions  of  substances  which  are  not  readily  volatile 
at  the  boiling-point  of  water.    Now,  we  have  in  the  first  part  of 
this  chapter  (112)  considered  the  relationship  between  boiling- 
points  and  vapor  pressures,  and  it  will  easily  be  understood  that 
a  solution  will  boil  at  the  temperature  at  which  the  pressure  of 
its  vapor  is  equal  to  the  atmospheric  pressure. 

129.  The  Lowering  of  the  Vapor  Pressure  by  Dissolved  Sub- 
stances.— If  a  solution  must  be  heated  above  100°  to  raise  its 
vapor  pressure  to  that  which  water  has  at  100°,  it  is  clear  that  at 
this  latter  temperature  the  solution  has  a  lower  vapor  pressure  than 
pure  water.    It  is  also  a  fact  that  at  every  lower  temperature  the 
vapor  pressure  of  a  solution  of  an  involatile  substance  is  less  than 
that  of  the  pure  solvent  at  the  same  temperature.    This  is  a 
very  important  universal  law.     The  law  applies  to  solutions 
formed  from  all  kinds  of  solvents. 

130.  Deliquescence. — In  the  case  of  a  very  soluble  substance, 
like  caustic  soda,  the  vapor  pressure  of  the  saturated  solution 
may  be  so  small  that  it  is  below  the  partial  pressure  exerted  by 
the  vapor  usually  present  in  the  air.     If  such  a  solution  is 
exposed  to  the  air,  water  vapor  from  the  air  will  condense  in  it 
until  the  solution  has  become  so  dilute  that  its  vapor  pressure 
is  just  equal  to  the  partial  pressure  of  the  water  vapor  in  the  air. 
Moreover,  if  such  a  very  soluble  substance  is  exposed  to  air  con- 
taining moisture,  water  will  condense  on  the  solid,  thus  convert- 
ing it  slowly,  first  into  a  saturated  solution,  and  finally  into  a 
dilute  solution.    This  action  is  called  deliquescence.    We  say 
caustic  soda  is  a  deliquescent  substance.    A  little  thought  will 
lead  to  the  conclusion  that  deliquescence  is  the  result  of  two  con- 
current conditions;  first,  the  possibility  of  the  formation,  by  a 
substance,  of  a  saturated  solution  which  has  a  very  small  vapor 
pressure  as  compared  with  pure  water — a  condition  usually 


78  Introduction  to  General  Chemistry 

accompanying  great  solubility;  and,  secondly,  the  presence  in  the 
air  of  a  sufficiently  great  water-vapor  content.  No  substance 
is  deliquescent  in  a  perfectly  dry  atmosphere,  while  every 
soluble  substance  exhibits  this  property  in  air  saturated  with 
water  vapor.  Deliquescence  is,  therefore,  not  a  fixed  property  of 
a  substance.  Thus  common  salt  is  usually  decidedly  deliquescent 
at  the  seashore,  where  the  air  contains  much  water  vapor;  but 
it  never  shows  this  property  in  a  desert  region. 

In  several  experiments  we  have  used  caustic  soda  or  calcium 
chloride  to  dry  air  or  other  gases  or  to  absorb  water  vapor  formed 
in  the  burning  of  hydrogen  (39,  50).  These  drying  agents  are 
among  the  most  deliquescent  substances  known. 

131.  Efflorescence. — In  paragraph  94  the  peculiar  behavior 
of  sodium  sulfate  decahydrate,  Na2SO4,  ioH20,  when  exposed 
to  the  open  air  was  described.     We  are  now  in  a  position  to 
understand  more  about  this  spontaneous  loss  of  water.     If  a 
crystal  of  the  hydrate  is  floated  on  the  surface  of  mercury  in  a 
vaccum  tube  like  one  of  those  shown  in  Figure  25,  the  mercury 
level  is  depressed  more  than  can  be  accounted  for  by  the  weight 
of  the  crystal.    Apparently  the  latter  is  giving  off  water  vapor 
and  attempting  to  establish  a  saturation  pressure.     This  pres- 
sure is  called  the  vapor  pressure  of  the  hydrate.    As  a  matter  of 
fact  all  hydrates  show  this  same  behavior,  with  the  difference 
that  each  has  its  own  characteristic  vapor  pressure  at  a  given 
temperature.    With  increased  temperature  the  vapor  pressure 
rises.    If  a  hydrate  is  exposed  to  air  in  which  the  partial  pressure 
of  water  vapor  is  less  than  the  vapor  pressure  of  this  substance, 
the  latter  will  give  off  water  to  the  air  just  as  a  water  surface 
does  to  air  in  which  the  partial  pressure  of  water  vapor  is  below 
the  saturation  value  for  water.    Along  with  the  loss  of  water, 
the  crystals  of  the  decomposing  hydrate  crumble  to  a  powder. 
This  process  is  called  efflorescence.    It  is  obvious  that  whether 
or  not  a  given  hydrate  effloresces  depends  not  only  upon  its  own 
vapor  pressure  but  upon  the  moisture  content  of  the  air  surround- 
ing it. 

132.  Effect  of  Temperature  on  Solubility.— The  solubility  of 
a  substance,  that  is,  the  amount  of  the  substance  which  dissolves 


Water  and  Solutions 


79 


(to  form  a  saturated  solution)  in  a  given  amount  of  water,  is 
dependent  upon  the  temperature.  Most  substances  are  more 
soluble  at  a  higher  than  at  a  lower  temperature;  but  this  is  not 
always  the  case,  as  the  solubility  of  some  substances  decreases 
with  rise  of  temperature.  In  fact,  gases  are  always  less  soluble 
at  a  higher  temperature. 


100 
90 
80 


I    70 


O 

.   60 
§» 

•H  O 

rtS   50 


40 


£ 

V< 

o 

03 

I20 


10 


d5> 


NaCl 


40         50        60 
Temperature 
FIG.  27 


70 


80 


90       TOO 


The  change  of  solubility  with  change  of  temperature  can 
most  easily  be  expressed  graphically,  that  is,  by  means  of 
so-called  solubility  curves.  The  accompanying  diagram  (Fig.  27) 
illustrates  the  method  and  gives  the  curves  for  water  solutions  of 
several  substances. 

133.  Effect  of  Crystalline  Form  on  Solubility. — Sodium  sul- 
fate  has  the  formula  NaaSCV  By  the  action  of  water  we  may 


80  Introduction  to  General  Chemistry 

readily  obtain  the  hydrate  NaaSCV  ioH2O  (96),  which  can  easily 
be  recrystallized  from  water,  as  described  under  "  Supersaturated 
Solutions"  in  this  chapter.  We  see  that  the  solubility  curve  for 
NaJSCV  ioH2O  rises  rapidly  until  a  temperature  of  33°  is  reached. 
At  this  temperature  the  crystals  melt  and  at  the  same  time 
decompose  into  Na2S04  and  H2O,  thus: 

Na,S(V  ioH2O  ->  Na2SO4+  ioH2O. 

Above  33°  we  have  the  solubility  curve  of  anhydrous  Na^SO^ 
which  is  a  different  chemical  substance  from  its  hydrate.  Thus 
we  see  that  there  are  for  the  anhydrous  salt  and  its  hydrate  two 
distinct  solubility  curves,  and  that  these  intersect  at  a  point 
for  which  the  temperature  is  that  at  which  the  hydrate  changes 
into  the  anhydrous  substance.  This  is  a  typical  case.  Each 
hydrate  of  a  substance  has  its  own  solubility  curve;  but  these  always 
intersect  at  the  point  corresponding  to  the  temperature  at  which  one 
substance  changes  into  the  other.  The  difference  in  solubility  is 
due  to  the  fact  that  each  has  its  own  characteristic  crystalline 
form. 

134.  Heat  of  Solution  and  Changes  of  Solubility  with 
Temperature. — A  question  which  will  now  very  naturally  occur 
to  the  student  is:  Why  should  the  solubility  of  various  sub- 
stances change  with  temperature  in  different  ways?  Although  a 
complete  and  satisfying  answer  cannot  be  given  to  this  question, 
it  is  possible  to  find  a  connection  between  the  shape  of  the  solu- 
bility curve  of  a  substance  and  another  fundamental  property. 
It  will  be  recalled  that  potassium  nitrate  absorbs  much  heat 
upon  dissolving  in  water,  and  we  notice  that  its  solubility  curve 
rises  rapidly  with  temperature.  Sodium  chloride  dissolves  with 
but  slight  absorption  of  heat,  and  its  curve  is  nearly  horizontal. 
Finally,  when  it  is  known  that  anhydrous  sodium  sulfate, 
Na2SO4,  dissolves  at  temperatures  above  33°  with  production  of 
heat,  and  that  its  curve  falls  with  rising  temperature,  the  general 
law  becomes  apparent.  These  are  typical  cases.  If  any  sub- 
stance dissolves  with  absorption  of  heat,  its  solubility  curve 
rises  with  rise  of  temperature.  If  it  dissolves  with  evolution  of 
heat,  then  the  curve  falls  with  rise  of  temperature.  The  frac- 


Water  and  Solutions  81 

tional  change  of  solubility  with  rise  of  i°  of  temperature  is  in 
general  proportional  to  the  heat  of  solution.  In  every  case  that 
change  of  solubility  which  will  absorb  heat  will  take  place  when  the 
temperature  is  raised.  This  will  involve  a  decrease  of  solubility 
with  rise  of  temperature,  in  the  case  of  a  substance  like  Na2S04, 
above  33°,  since,  if  heat  is  evolved  when  the  substance  dissolves, 
heat  is  absorbed  in  equal  amount  when  the  same  weight  of  the 
substance  crystallizes  out  of  a  solution. 

In  some  cases  where  heat  is  evolved  when  a  substance  is 
dissolved,  the  observed  heat  is  the  result  of  the  union  of  the  solid 
with  water  to  form  a  hydrate,  which  may  dissolve  with  a  small 
absorption  of  heat.  In  such  cases  the  solubility  of  the  hydrate 
increases  with  rise  of  temperature  in  strict  accord  with  the  law. 
For  example  when  anhydrous  calcium  chloride  is  dissolved  in 
water  the  mixture  gets  very  hot.  The  saturated  solution 
deposits  crystals  of  CaCl2,  6H20  on  cooling.  This  hydrate  dis- 
solves in  water  with  absorption  of  heat  and  its  solubility  increases 
with  a  rise  in  temperature.  The  heat  given  out  on  dissolving  the 
anhydrous  salt  is  the  excess  of  the  heat  produced  in  the  reaction 

CaCl2+6H20  =  CaCl2,6H2O 

above  the  heat  absorbed  in  the  dissolving  of  the  hydrate  CaCl2, 
6H20. 

135.  Two  Apparent  Kinds  of  Solubility. — In  cases  of  ordinary 
solubility,  evaporation  of  the  water  leads  to  the  recovery  un- 
changed of  the  substance  originally  dissolved.  In  other  cases, 
evaporation  of  the  solution  obtained  by  the  apparent  dissolving 
of  a  substance  leaves  an  entirely  different  substance.  For 
example,  if  we  throw  a  piece  of  sodium  on  water  the  former  soon 
disappears  and  a  solution  results  (40, 88) .  We  might  be  inclined 
to  say  that  the  sodium  has  dissolved  in  the  water;  but  there  is 
another  way  of  looking  at  the  matter.  We  know  that  in  this 
case  a  chemical  change  has  occurred,  as  represented  by  the 
equation 

2Na+  2H2O  ->  2NaOH+H2. 

Furthermore,  we  know  that  by  evaporation  of  the  solution  we 
get  sodium  hydroxide  and  not  sodium;  for  this  reason  it  seems 


82  Introduction  to  General  Chemistry 

more  logical  to  say  that  sodium  and  water  react  to  give  sodium 
hydroxide,  which  then  dissolves  in  water,  than  to  say  that 
sodium  itself  is  soluble  in  water.  In  fact,  we  know  nothing 
about  the  solubility  of  sodium  in  water,  since  the  two  react  as 
soon  as  they  are  brought  into  contact.  We  know  a  very  great 
number  of  cases  analogous  to  this  one,  and  in  all  of  them  we 
recognize  that  we  have  to  deal  with  chemical  changes  which 
give  rise  to  soluble  products. 

136.  Normal  Solutions. — In  the  neutralization  of  hydro- 
chloric acid  by  sodium  hydroxide,  which  takes  place  according 
to  the  equation 

HCl+NaOH  ->  NaCl+H20,  (89) 

one  formula  weight  of  the  acid  (36. 5  g.)  requires  one  formula 
weight  of  the  base  (40  g.).  If  we  make  a  solution  of  the  acid  of 
such  concentration  that  i  liter  contains  36-5g.  of  hydrogen 
chloride,  and  also  make  a  solution  of  the  base  containing  40  g. 
of  sodium  hydroxide  per  liter,  then  upon  mixing  the  liter  of  the 
acid  solution  with  the  liter  of  the  basic  solution  exact  neutraliza- 
tion will  take  place.  It  follows,  of  course,  that,  to  neutralize  a 
given  volume  of  such  an  acid  solution,  exactly  the  same  volume 
of  the  basic  solution  will  be  required.  We  call  such  solutions 
normal  solutions. 

If  we  wish  to  make  a  solution  of  nitric  acid  of  such  concentra- 
tion that  i  liter  of  it  will  exactly  neutralize  i  liter  of  normal 
sodium  hydroxide,  we  see,  in  accord  with  the  equation 

HNO3+NaOH  ->  NaN03+H2O,  (104) 

that  one  formula  weight  of  HN03  must  be  contained  in  i  liter 
of  the  solution.  This  gives  a  normal  solution  of  nitric  acid. 

Now  the  case  is  a  little  different  if  a  normal  solution  of  sul- 
furic  acid  is  to  be  made,  since  in  this  case  we  have 

H2SO4+  2NaOH  ->  Na2SO4+  2H2O.  (94) 

We  see  that  one  formula  weight  of  sulfuric  acid  neutralizes  two 
formula  weights  of  sodium  hydroxide,  so  that  to  neutralize 
i  liter  of  normal  sodium  hydroxide,  which  contains  but  one 
formula  weight  of  the  base,  only  one-half  a  formula  weight 


Water  and  Solutions 


FIG.  28 


(|  of  98  g.  or  49  g.)  of  sulfuric  acid  is  required.  Therefore  if  we 
dissolve  49  g.  of  the  acid  in  sufficient  water  to  make  a  liter  of 
solution,  this  liter  of  acid  solution  will  just  neutralize  i  liter  of 
normal  sodium  hydroxide.  We  call  the  sulfuric 
acid  solution  so  made  also  a  normal  solution. 

A  normal  solution  of  potassium  hydroxide, 
KOH,  would  contain  one  formula  weight  (56  g.) 
per  liter  (106).  A  normal  solution  of  any  acid 
always  neutralizes  an  exactly  equal  volume  of  a 
normal  solution  of  any  base.  The  term  "normal" 
is  usually  abbreviated  N,  so  that  for  a  normal 
solution  by  hydrochloric  acid  we  write  N  HC1. 
Normal  solutions  are  of  great  importance  in  prac- 
tical work.  Suppose  we  wish  to  know  the  concen- 
tration of  a  given  solution  of  sodium  hydroxide. 
We  take,  with  a  pipette  (Fig.  28),  a  carefully 
measured  volume,  say  20  c.c.,  add  to  it  sufficient 
litmus  solution  to  produce  a  pale  blue  color,  and  then  from  a 
measuring  tube,  called  a  burette  (Fig.  29),  run  in  a  normal  solu- 
tion of  hydrochloric  or  other  acid  until  the  color  just  changes 
from  blue  to  red.  A  little  practice  enables  one  to  find,  to  within 

one  drop  or  less,  the  volume  of  acid 
required.  Let  us  say  42  c.c.  of 
N  HC1  was  required  for  the  20  c.c. 
of  NaOH  solution  of  unknown  con- 
centration. Our  problem  is  to  find 
the  weight  of  sodium  hydroxide  in 
the  20  c.c.  of  solution  taken.  Now, 
42  c.c.  of  N  acid  will  neutralize 
42  c.c.  of  N  sodium  hydroxide,  of 
which  i  liter  (  =  1,000  c.c.)  contains 
40  g.  of  sodium  hydroxide.  There- 
fore the  weight  of  sodium  hydroxide  in  the  20  c.c.  taken  =  o .  042  X 
40  g.  =  i .  68  g.  We  also  see  that  the  sodium  hydroxide  solution 
is  42/20  =  2.1  times  as  concentrated  as  a  normal  solution  of  this 
base.  We  express  its  concentration  by  saying  that  it  is  2.1 
times  normal  in  concentration, 


FIG.  29 


84  Introduction  to  General  Chemistry 

It  is  often  convenient  in  practice  to  use  solutions  of  £,  £,  TV , 
or  some  other  fraction  of  normal;    we  call  these  half-normal 

/N\  /N\  /N\ 

(  —  ),  one-fifth  normal  I  —  ),  and  one-tenth  or  deci-normal  (  —  ), 

respectively. 

137.  Acidimetry  and  Alkalimetry. — The  analyses  of  acids 
and  bases  by  means  of  normal  solutions  are  called  respectively 
acidimetry  and  alkalimetry.    The  act  of  running  in  a  solution 
from  a  burette  until  the  neutral  or  end-point  is  reached  is  called 
titration.    The  volume  of  solution  used  is  called  the  titer. 
Instead  of  litmus  we  may  use  some  other  colored  substance  to 
indicate  the  end-point;  such  a  substance  is  called  an  indicator. 
Other  useful  indicators  are  methyl  orange,  phenolphthalein, 
and  Congo  red. 

138.  Problems. — 

N 

1.  How  many  c.c.  of  —  nitric  acid  are  required  to  neutralize 

50  c.c.  of  normal  potassium  hydroxide?     (107) 

N 

2.  How  many  c.c.  of  —  sodium  hydroxide  are  required  to 

o 

N 
neutralize  20  c.c.  of  —  sulfuric  acid?     (94) 

0 

3.  If  1 6  c.c.  of  a  solution  of  sulfuric  acid  of  unknown  concen- 

N 
tration  requires  for  its  neutralization  36  c.c.  of  —  potassium 

hydroxide,  (a)  what  is  the  weight  of  sulfuric  acid  in  the  16  c.c. 
taken?  (b)  what  is  the  weight  of  sulfuric  acid  in  i  liter  of  this 
acid?  (107) 

139.  The  Formation  of  Water. — We  have,  in  earlier  chapters, 
learned  various  ways  in  which  water  can  be  formed  chemically. 
We  may  enumerate  these  by  way  of  review. 

Water  is  formed — 

1.  By  the  burning  of  hydrogen: 

2H2+O2->2H20. 

2.  By  the  burning  of  a  compound  of  hydrogen,  for  example, 

methane: 

CH4+2Oa->CO2+2HaO.  (86) 


Water  and  Solutions  85 

3.  By  the  oxidation  of  hydrogen  or  its  compounds  by  means 
of  combined  oxygen,  as,  for  example,  when  ammonia  is  passed 
over  hot  copper  oxide : 

2NH3+3CuO  ->  3H2O+3Cu+N2.  (84) 

4.  By  the  union  of  acids  and  bases,  whereby  a  salt  and  water 
are  always  formed;  for  example: 

HCl+NaOH  ->  NaCl+H20.  (89) 

5.  By  the  decomposition  or  dissociation  of  various  unstable 
compounds,  as,  for  example,  sodium  sulfate  decahydrate  into 
the  anhydrous  salt  and  water: 

Na2SO4-  ioH2O  ->  NaSO4+  ioH2O.  (96) 

140.  The  Chemical  Reactions  of  Water. — We  have  also 
studied  some  of  the  important  kinds  of  reactions  in  which  water 
takes  part.  We  may  now  summarize  these  as  follows: 

1 .  Water  unites  with  salts  to  form  hydrates,  thus : 

Na2SO4+  ioH20  ->  Na2S04-  ioH2O.  (96) 

2.  Ammonia  and  water  unite  to  form  ammonium  hydroxide: 

NH3+H2O  ->  NH4OH.  (91) 

4.  Water  acts  upon  some  metals  to  give  hydroxides  and 
hydrogen.  Thus,  sodium  and  cold  water  react  very  easily, 
giving  sodium  hydroxide  and  hydrogen: 

2Na+  2H2O  ->  2NaOH+H2.  (88) 

Magnesium  does  not  act  readily  on  cold  water,  but  burns  vigor- 
ously in  steam  giving  the  hydroxide  and  hydrogen : 

Mg+  2H20  ->  Mg(OH)2+H2.  (28,  86) 


CHAPTER  IX 
ACIDS,  BASES,  AND  SALTS.— II 

141.  New  Acids,  Bases,  and  Salts. — The  present  chapter  will 
treat  of  three  new  acids,  carbonic,  H2C03,  acetic,  C2H4O2,  and 
phosphoric,  H3P04,  and  the  bases  derived  from  the  elements 
magnesium,   calcium,   barium,   zinc,   iron,   aluminum,   copper, 
silver,  lead,  and  mercury,  together  with  the  more  important  salts 
which  these  bases  form  with  the  three  acids  studied  in  the  first 
chapter  on  acids,  bases,  and  salts,  as  well  as  with  the  three  acids 
above-mentioned. 

142.  The  Action  of  Water  on  Magnesium  Oxide :  Magnesium 
Hydroxide,  Mg(OH)2. — All  of  the  three  bases  studied  in  the  first 
chapter  on  "Acids,  Bases,  and  Salts"  are  readily  soluble  in  water. 
We  shall  next  consider  one  which  dissolves  in  water  only  to  a  very 
slight  extent.     If  we  shake,  with  water,  a  little  magnesium 
oxide  (u,  80),  obtained  by  burning  magnesium,  we  find  that  the 
solution  will  turn  red  litmus  blue,  although  but  a  small  amount 
of  the  magnesium  oxide  has  dissolved  in  the  water,  the  larger 
part  having  remained  undissolved.    It  has  been  found  by  careful 
experiment  that  magnesium  oxide  and  water  unite  when  brought 
together,  giving  a  single  new  •  compound,  the  composition  of 
which  is  represented  by  Mg02H2,  which  we  may  also  write 
Mg(OH)2,  and  call  magnesium  hydroxide.    This  is  a  white  sub- 
stance, with  which  the  student  may  already  be  familiar  under  the 
name  of  " milk  of  magnesia."    It  is  extensively  used  in  medicine. 
It  is  to  be  classified  as  a  base,  since,  like  sodium  hydroxide,  it 
colors  litmus  blue  and  neutralizes  hydrocloric  acid.     The  equa- 
tion for  the  action  of  water  on  the  oxide  is 

MgO+H2O-»Mg(OH)2. 

143.  The  Action  of  Hydrochloric  Acid  on  Magnesium  Hy- 
droxide:  Magnesium  Chloride,  MgCl2.— Magnesium  hydroxide 
is  but  very  slightly  soluble  in  water.    However,  if  we  add  hydro- 
chloric acid  to  the  magnesium  hydroxide  formed  from  the  mag- 

86 


Acids,  Bases,  and  Salts — //  87 

nesium  oxide  and  water  until  the  solution  just  turns  litmus  red, 
we  find  that  all  of  the  solid  dissolves,  giving  a  clear,  colorless 
solution  which  if  left  to  evaporate  in  an  open  vessel  will  deposit 
colorless  crystals.  An  investigation  of  this  new  substance  shows 
that  it  is  a  compound  of  magnesium,  chlorine,  hydrogen,  and 
oxygen,  in  the  proportion  indicated  by  the  formula  MgCl2-  6H2O. 
This  hydrate  of  magnesium  chloride  is  formed  as  a  result  of  the 
following  two  reactions: 

Mg(OH)2+  2HC1  ->  MgCl2+  2H2O, 
MgCl2+6H2O  ->  MgCL- 6H2O. 

144.  Magnesium  Sulfate,  MgSO4. — If  we  now  add  diluted 
sulfuric  acid  to  some  magnesium  hydroxide  mixed  with  water, 

1  until  all  of  the  solid  has  dissolved  and  litmus  shows  the  solution 
to  be  neutral,  we  may  obtain  from  the  solution  by  careful 
evaporation  crystals  of  magnesium  sulfate  having  the  formula 
MgSCV  7H20,  a  substance  much  used  in  medicine  and  known  as 
Epsom  salts.  The  reaction  occurs  according  to  the  equation : 

H2S04+Mg(OH)2  ->  MgS04+  2H20, 

the  MgSO4  then  combining  with  water  from  the  solution,  thus: 
MgS04+7H20-»  MgS04..7H20, 

to  form  the  hydrate.  The  latter,  when  heated,  readily  disso- 
ciates into  MgSO4  and  7H2O,  a  fact  which  may  be  expressed  thus: 

MgSO4  •  7H2O  ->  MgSO4+  7H2O. 

This,  as  we  see,  is  just  the  reverse  of  the  preceding  reaction. 
The  reactions  of  hydrates  in  solution  are  of  course  the  same  as 
those  of  the  anhydrous  salts,  since  solutions  of  the  two  cannot 
be  distinguished.  In  what  follows,  the  discussion  of  the  hydrates 
formed  will  be  omitted  except  in  a  single  important  instance. 

145.  Magnesium  Nitrate,  Mg(NO3)2. — Magnesium  hydroxide 
is  readily  neutralized  by  nitric  acid,  with  the  formation  of 
magnesium  nitrate,  which  forms  white  crystals  very  easily 
soluble  in  water: 

Mg(OH)2+  2HN03  ->  Mg(N03)2+  2H2O. 


88  Introduction  to  General  Chemistry 

Magnesium  oxide  and  dilute  hydrochloric  acid  react  to  give 
magnesium  chloride,  which  is  the  same  compound  as  that  formed 
from  magnesium  hydroxide  and  the  same  acid.  The  equation 
for  the  reaction  is 

MgO+  2HC1  ->  MgCl2+H2O. 

The  corresponding  reactions  take  place  with  sulfuric  and  with 
nitric  acid,  and  are  represented  by  the  equations 

MgO+H2SO4  ->  MgSO4+H2O, 
MgO+  2HNO3  ->  Mg(NO3)2+H2O. 

146.  Monacid  and  Diacid  Bases:  Valence. — If  we  compare 
the  formula  of  magnesium  chloride,  MgCl2,  with  that  of  sodium 
chloride,  NaCl,  or  potassium  chloride,  KC1,  we  see  that,  in  the 
first  case,  one  symbol  weight  of  the  metal  is  combined  with  two 
symbol  weights  of  chlorine,  while  in  the  other  two  cases  one 
symbol  weight  of  metal  is  combined  with  but  one  symbol  weight 
of  chlorine.     In  the  cases  of  the  neutralization  of  the  hydroxides 
by  hydrochloric  acid  we  found  that  one  formula  weight  of  mag- 
nesium hydroxide  required  two  formula  weights  of  hydrochloric 
acid  (143) ;  while  one  formula  weight  of  the  hydroxide  of  either 
sodium  or  potassium  required  but  one  of  hydrochloric  acid  (102, 
107).     For  this  reason  we  call  sodium  and  potassium  hydroxides 
monacid  bases,  and  magnesium  hydroxide  a  diacid  base.    We 
also  make  use  of  the  term  valence  in  referring  to  facts  like  those 
just  mentioned,  saying  that  the  valence  of  sodium  or  potassium 
is  one,  while  that  of  magnesium  is  two,  or  that  sodium  and 
potassium  are  univalent,  while  magnesium  is  a  bivalent  element. 
Since  hydrogen  chloride  has  the  formula  HC1,  we  say  that  hydro- 
gen has  a  valence  of  one,  and  we  also  say  that  the  valence  of 
chlorine  is  one. 

147.  Radicals  and  Their  Valence. — We  have  already  become 
acquainted  with  several  ammonium  compounds,  as,  for  example, 
the  chloride  NH4C1  (91)  and  the  nitrate  NH4NO3  (105).    We 
call  the  combination  NH4  the  ammonium  radical;  it  has  never 
been  obtained  as  a  separate  substance,  but  is  known  only  as  a 
component  of  ammonium  compounds.    We  know  of  many  other 
such  radicals,  one  of  which  is  met  with  in  sulfuric  acid  and  sul- 


Acids,  Bases,  and  Salts — II  89 

fates,  where  we  have  found  that  sulfur  and  oxygen  are  always 
present  in  the  ratio  represented  by  SO4.  Here  again  we  have  a 
radical  which  is  found  in  many  salts,  the  sulfates,  but  is  not 
known  as  a  separate  chemical  substance.  In  nitric  acid  and 
the  nitrates  we  have  the  radical  N03.  A  radical  is  composed  of 
two  or  more  elements  united  in  a  definite  proportion;  con- 
sequently the  composition  of  a  radical  can  always  be  represented 
by  a  formula.  We  may  consider  that  the  combination  of  nitro- 
gen and  hydrogen,  NH4,  taken  as  a  radical,  has  a  valence  of  one, 
since  the  weight  of  nitrogen  and  hydrogen  represented  by  NH4 
taken  once  unites  with  one  symbol  weight  of  chlorine,  giving 
NH4C1.  Since  sulfuric  acid,  H2S04,  forms  such  salts  as  Na2S04, 
K2SO4  and  (N'H4)2SO4,  we  say  that  the  sulfate  radical,  S04,  has 
a  valence  of  two,  a  fact  which  is  also  shown  by  the  existence  of 
such  salts  as  NaHSO4,  etc.  Now  if  magnesium  has  a  valence 
of  two  and  the  sulfate  radical  has  also  the  valence  of  two,  we 
see  in  the  fact  that  magnesium  sulfate  has  the  formula  MgSO4,  a 
broader  meaning  of  the  term  valence.  And  so  chemists  often 
speak  of  the  two  valences  of  magnesium  being  satisfied  by  the 
two  valences  of  the  sulfate  radical.  The  subject  of  valence  will 
be  considered  again  at  the  end  of  this  chapter  (183). 

148.  Zinc  and  Its  Salts. — Zinc  is  an  element  and  a  very 
important  metal;  it  is  known  in  commerce  as  spelter,  and  is  used 
in  enormous  amounts  in  making  galvanized  iron,  which  is  iron 
coated  with  metallic  zinc,  in  making  brass,  whose  other  com- 
ponent is  copper,  and  for  many  other  purposes.  Zinc  will  burn 
when  strongly  heated  in  the  air  or  in  oxygen,  giving  a  white 
oxide,  the  reaction  being  represented  by  the  equation 

2Zn+O2->2ZnO. 

Zinc  oxide  is  used  extensively  in  making  white  paint. 

It  will  be  recalled  that  magnesium  burns,  giving  an  oxide 
(n,  80),  and  that  this  oxide  reacts  with  acids  giving  salts,  thus: 

MgO+2HCl->MgCl2+H20.  (143,  MS) 

Zinc  oxide  behaves  like  magnesium  oxide  when  treated  with 
hydrochloric  acid,  giving  zinc  chloride,  thus: 
ZnO+2HCl->  ZnCl2+H2O. 


go  Introduction  to  General  Chemistry 

Zinc  chloride  is  a  salt  which  dissolves  very  readily  in  water, 
giving  a  clear,  colorless  solution.  Zinc  oxide  gives  zinc  sulfate 
and  zinc  nitrate  as  follows: 

ZnO+H2SO4->  ZnSO4+H2O, 
ZnO+  2HNO3  ->  Zn(NO3)2+H2O. 

These  are  white  salts,  also  easily  soluble  in  water. 

149.  The  Action  of  Hydrochloric  Acid  on  Zinc. — If  we  pour 
some  hydrochloric  acid  on  zinc  we  observe  a  vigorous  reaction; 
the  zinc  dissolves  and  a  gas  which  proves  to  be  hydrogen  is  given 
off.     If,  after  the  zinc  has  all  dissolved,  we  evaporate  the  solution, 
we  obtain  a  white  solid  which  is  found  to  be  zinc  chloride.     The 
reaction  is  represented  thus: 

Zn+2HCl->ZnCl2+H2. 

Comparing  this  equation  with  that  for  the  action  of  zinc  oxide 
on  hydrochloric  acid,  we  see  that  in  the  latter  case  the  hydrogen 
of  the  acid,  instead  of  passing  off  as  gas,  unites  with  the  oxygen 
of  the  zinc  oxide,  giving  water. 

We  might  expect  that  metallic  magnesium  and  hydrochloric 
acid  would  act  thus : 

Mg+2HCl->MgCl2+H2, 

and  it  is  easy  to  show  by  experiment  that  this  is  the  case.  With 
dilute  sulf uric  acid  these  metals  behave  as  follows : 

Mg+H2S04->MgS04+H2, 
Zn+H2SO4  ->  ZnSO4+H2. 

In  making  hydrogen  in  the  laboratory  we  usually  use  zinc  and 
hydrochloric  acid. 

150.  Marble  and  Other  Compounds  of  the  Element  Cal- 
cium.— Let  us  now  consider  the  chemical  behavior  of  marble. 
If  we  place  some  lumps  of  marble  in  a  hard  glass  tube  and  heat 
strongly,  a  gas  is  given  off,  while  the  lumps  change  but  little  in 
appearance.    This  gas  causes  limewater  to  turn  milky;    it  is 
carbon  dioxide  (19).     If  the  lumps  left  after  heating  the  marble 
are  moistened  with  water,  they  grow  very  hot,  swell  up,  and 
crumble  to  a  white  powder.     It  is  evident  therefore  that  the 


Acids,  Bases,  and  Salts — II  91 

marble  has  been  changed  chemically  by  the  heating.  The  solid 
left  after  the  heating  is  the  common  substance,  quicklime.  The 
action  of  water  upon  quicklime  is  called  slaking.  If  the  slaked 
lime  is  shaken  with  a  large  amount  of  water,  not  much  seems  to 
dissolve;  but  if  we  filter  the  mixture,  a  clear,  colorless  solution 
is  obtained.  If  some  carbon  dioxide  gas  is  run  into  this  clear 
solution  it  turns  milky,  because  this  solution  is  limewater  (18), 
of  which  we  have  so  often  made  use.  If  we  test  the  limewater 
with  litmus  we  find  that  it  turns  the  latter  blue,  showing  the 
limewater  to  be  a  solution  of  a  base.  This  base  reacts  with  acids 
to  form  salts.  All  of  these  products  contain  an  element  called 
calcium,  whose  symbol  is  Ca.  Calcium  is  a  brassy-looking  metal, 
which  will  readily  burn  with  a  bright  light  if  heated  in  air  or 
oxygen,  giving' calcium  oxide: 

2Ca-f-O2->2CaO. 

Calcium  oxide,  CaO,  is  quicklime;  but  the  latter  is  never 
made  practically  in  this  way,  because  metallic  calcium  is  too 
expensive,  and  because  the  oxide  is  made  very  cheaply  by  heating 
marble  or,  more  often,  limestone,  which  is  an  impure  form  of  the 
same  compound  as  marble.  By  heating  marble,  CaO  and  CO2 
are  formed,  and  nothing  else.  By  finding  the  percentage  of 
each  we  can  easily  calculate  the  formula  for  marble  to  be  CaC03, 
which  is  called  calcium  carbonate;  the  effect  of  the  heating  is, 
therefore,  represented  thus:  . 

CaCO3->CaO+CO2. 

151.  Calcium  Hydroxide,  Ca(OH)2. — As  has  been  stated, 
when  water  acts  on  calcium  oxide  or  quicklime,  we  get  calcium 
hydroxide  or  slaked  lime,  a  solution  of  which  is  called  limewater : 

CaO+H2O->Ca(OH)2. 

This  reaction  is  analogous  to  the  action  of  water  on  magnesium 
oxide,  which  was  studied  earlier  (142).  The  action  of  hydro- 
chloric acid  on  calcium  hydroxide  gives  calcium  chloride  and 
water : 

Ca(OH)2+2HCl  ->  CaCl2+2H2O. 


Q2  Introduction  to  General  Chemistry 

We  are  now  in  position  to  understand  the  cause  of  the  milki- 
ness  produced  when  carbon  dioxide  acts  on  limewater.  The 
white  solid  formed  is  really  calcium  carbonate,  CaCO3.  The 
equation  is 

Ca(OH)2+CO2  ->  CaCO3+H2O. 

152.  Carbonic  Acid,  H2CO3. — If  we  pass  carbon  dioxide  into 
water,  a  solution  results  which  has  faint  acid  properties.     This 
solution  is  in  fact  the  well-known  plain  soda  served  at  soda 
fountains.    The  dissolved  carbon  dioxide  and  water  partially 
combine  to  form  an  acid  called  carbonic  acid : 

CO2+H2O->H2CO3. 

Therefore  we  may  then  consider  that  it  is  this  acid  which 
neutralizes  the  base  calcium  hydroxide,  thus: 

Ca(OH)2+H2C03->  CaCO3+2H2O. 

Calcium  carbonate  is  a  salt  which  is  almost  insoluble  in  water. 
In  fact,  salts  exhibit  all  degrees  of  solubility  in  water.  Some, 
like  zinc  chloride,  dissolve  in  less  than  their  own  weight  of  water; 
others,  like  common  salt,  are  much  less  soluble;  while  many,  like 
calcium  carbonate,  are  very  nearly  insoluble  in  water. 

153.  Calcium  Sulfate,  CaSO4. — Calcium  hydroxide  and  sul- 
furic  acid  form  calcium  sulfate  and  water: 

Ca(OH)2+H2S04->  CaSO4+2H2O. 

We  find  by  experiment  that  the  calcium  sulfate  so  formed  dis- 
solves very  slightly  in  water,  100  c.c.  of  water  dissolving  but 
one-fourth  of  a  gram  of  the  salt.  On  the  other  hand,  calcium 
chloride  is  very  soluble  in  water.  If  we  add  to  a  solution  con- 
taining, say,  5  or  10  per  cent  of  calcium  chloride,  a  sufficient 
amount  of  sulfuric  acid,  we  observe  that  a  large  amount  of  a 
white  powder  forms  in  the  mixed  solutions  and  soon  settles  to 
the  bottom  of  the  vessel,  leaving  a  clear,  colorless  liquid  above. 
The  white  powder  proves  to  be  calcium  sulfate,  which  is  formed 

thus: 

CaCl2+H2SO4->  CaSO4+2HCl. 


Acids,  Bases j  and  Salts — //  93 

154.  Precipitation. — We  often  encounter  chemical  reactions 
in  which,  as  in  the  action  between  calcium  chloride  and  sulfuric 
acid,  a  solid  is  formed  upon  bringing  together  two  solutions.    A 
solid  so  thrown  down  is  called  a  precipitate ;  and  we  speak  of  the 
precipitation  of  calcium  sulfate.    The  formation  of  insoluble 
calcium  carbonate  by  the  action  of  carbon  dioxide  on  lim^vater 
is  another  example  of  precipitation. 

155.  Gypsum  and  Plaster  of  Paris. — Calcium  sulfate  occurs 
in  nature  as  the  mineral  gypsum,  CaSO4'2H2O.     If  part  of  the 
water  of  hydration  is  driven  off  by  heat,  gypsum  is  converted 
into  the  well-known  plaster  of  Paris: 

2CaS04-2H2O->  2CaSO4-H2O+3H2O. 

When  powdered  plaster  of  Paris  is  mixed  with  enough  water  to 
form  a  paste,  it  sets  in  the  course  of  an  hour  into  a  solid  mass 
which  retains  the  form  of  the  vessel  or  mold  which  holds  it. 
Plaster  casts  are  made  in  this  way.  The  setting  is  due  to  the 
formation  of  interlacing  crystal  filaments  of  the  hydrate 
CaSO4-  2H20,  formed  by  a  reversal  of  the  action  by  which  plaster 
of  Paris  is  formed  from  gypsum: 

2CaS04-H2O+3H2O->2CaSO4-2H20. 

156.  Calcium  Bicarbonate  and  Hard  Water. — A  very  inter- 
esting and  important  reaction  occurs  when  carbon  dioxide  is 
passed  for  a  long  time  into  a  sufficiently  dilute  solution  of  calcium 
hydroxide  (limewater) .    At  first  a  milkiness  appears,  due  to  the 
formation  of  calcium  carbonate : 

Ca(OH)2+CO2->  CaC03+H2O. 

If  we  continue  to  pass  in  carbon  dioxide,  the  precipitate  slowly 
dissolves,  giving  finally  a  perfectly  clear  solution.  If  this  solu- 
tion is  now  boiled,  carbon  dioxide  gas  is  given  off  and  a  white 
precipitate  is  formed.  These  facts  are  explained  in  the  following 
way.  Carbonic  acid,  H2C03,  like  sulfuric  acid,  is  a  dibasic 
acid  (102)  and  can  form  acid  salts  as  well  as  neutral  salts.  Just 
as  sulfuric  acid  yields  Na2S04  and  NaHSO4,  so  carbonic  acid  gives 
Na2C03  and  NaHC03,  sodium  carbonate  and  sodium  acid  car- 
bonate, also  known  as  bicarbonate  (baking-soda). 


94  Introduction  to  General  Chemistry 

The  calcium  salts  corresponding  to  sodium  carbonate  and 
bicarbonate  are  CaCO3  and  Ca(HC03)2.  The  difference  in  the 
formulae  of  the  sodium  and  calcium  salts  is  due  to  the  fact  that 
the  valence  of  calcium  is  two,  while  that  of  sodium  is  one.  Now 
when  carbon  dioxide,  in  excess,  acts  on  calcium  carbonate, 
calcium  acid  carbonate,  called  also  bicarbonate,  is  formed,  and 
this  being  soluble  in  water  the  precipitate  goes  into  solution : 

CO2+H2O-»H2C03 
CaCO3+H2CO3-»  Ca(HCO3)2. 

When  the  clear  solution  so  obtained  is  boiled  the  following  reac- 
tion occurs: 

Ca(HCO3)2->  CaCO3-J-  H2O+CO2. 

These  reactions  take  place  extensively  in  nature.  Natural 
waters,  e.g.,  those  of  springs  and  rivers,  contain  dissolved  carbon 
dioxide,  and  therefore  carbonic  acid.  Such  waters  passing  over 
limestone,  impure  CaCO3,  dissolve  it  and  take  the  Ca(HC03)2 
into  solution,  forming  so-called  hard  water.  When  boiled,  as 
in  a  teakettle,  it  gives  off  carbon  dioxide  and  deposits  the  calcium 
carbonate. 

157.  Vinegar:    Acetic   Acid,    C2H4O2. — Acetic    acid   is   the 
principal  ingredient,  other  than  water,  in  vinegar,  of  which  it 
constitutes  about  4  per  cent.    The  formula  of  acetic  acid  is 
C2H4O2.    It  neutralizes   sodium  hydroxide   according  to   the 

equation 

C2H402+NaOH=NaC2H302+H20. 

The  salt  NaC2H302,  sodium  acetate,  is  the  only  sodium  salt 
which  can  be  made  from  this  acid.  Therefore  the  acid  radical 
of  acetic  acid  and  its  salts  is  C2H302  and  we  may  write  the 
formula  of  the  acid  HC2H3O2  to  indicate  that  only  one  of  the  four 
hydrogen  atoms  of  a  molecule  is  replaceable  in  salt  formation. 

Pure  acetic  acid  is  a  colorless  liquid,  miscible  with  water  in 
all  proportions.  It  is  monobasic  and  forms  with  most  bases  salts 
called  acetates. 

158.  Bone    Ash:     Calcium    Phosphate,    Ca3(PO4)2.— When 
bones  are  burned  only  the  gelatinous  matter  and  connective 
tissue  are  removed;   the  white  material  which  is  left  is  called 


Acids,  Bases,  and  Salts — //  95 

bone  ash  and  consists  essentially  of  calcium  phosphate,  Ca3(P04)2. 
If  powdered  bone  ash,  which  is  practically  insoluble  in  water,  is 
stirred  with  somewhat  diluted  sulfuric  acid,  the  following  reac- 
tion occurs: 

Ca3(P04)2+3H2S04  ->  3CaS04+  2H3PO4. 

The  calcium  sulf ate  formed  is  difficultly  soluble  in  water  and  may 
be  filtered  out,  giving  a  clear,  colorless  filtrate  containing  dis- 
solved phosphoric  acid,  H3P04. 

159.  Phosphoric  Acid :  a  Tribasic  Acid. — This  acid  is  a  white 
crystalline  solid,  which  is  very  soluble  in  water,  frequently  com- 
ing on  the  market  in  the  form  of  a  very  concentrated  solution  of 
syrupy  consistency.    Its  dilute  solution  has  a  pleasant  sour  taste 
and  turns  litmus  red.    With  suitable  proportions  of  sodium 
hydroxide  it  yields  the  three  salts,  Na3P04,  trisodium  phosphate, 
Na2HPO4,    disodium    hydrogen     phosphate,     and    NaH2P04, 
sodium  dihydrogen  phosphate.    The  latter  is  a  typical  acid  salt, 
having  a  sour  taste  and  acid  action  on  litmus.    Phosphoric  acid 
is  therefore  a  tribasic  acid.    //  forms  with  bases  three  series  o] 
salts,  corresponding  to  those  of  sodium.    To  distinguish  these 
classes  of  salts  from  one  another  they  are  called  primary,  second- 
ary, and  tertiary,  that  with  the  smallest  proportion  of  base  being 
the  primary  and  that  in  which  all  hydrogen  is  replaced  being  the 
tertiary. 

1 60.  The   Practical  Importance   of   Calcium  Phosphate. — 
Since  calcium  phosphate,  Ca3(PO4)2,  constitutes  the  mineral 
matter  of  bones,  it  is,  of  course,  a  substance  of  very  great  im- 
portance.   Phosphates  in  small  amounts  are  also  indispensable 
constituents  of  most  plants,  and  it  is  from  these,  especially  from 
the  seeds,  like  wheat,  oats,  and  corn,  that  men  and  animals  get 
their  needed  supply.    Plants,  in  turn,  get  their  phosphates  from 
the  soil,  and  do  not  thrive  on  soil  deficient  in  phosphates.    Such 
infertile  soil  may  be  greatly  improved  by  the  use  of  fertilizers 
containing  phosphates.    For  this  purpose,  bone  ash  is  often 
employed;  but  since  bone  ash  is  almost  insoluble  in  water,  it  is 
not  directly  available  for  plant  use.    In  order  to  make  it  available 
it  is  treated  with  sufficient  sulfuric  acid  to  convert  it  into 


g6  Introduction  to  General  Chemistry 

Ca(H2PO4)2,  usually  known  as  calcium  superphosphate,  which 
is  soluble  in  water: 

Ca3(PO4)2+  2H2SO4  ->  Ca(H2PO4)2+  2CaSO4. 

Immense  deposits  of  calcium  phosphate  occur  in  Florida  and 
Tennessee,  as  phosphate  rock.  These  deposits  have  doubtless 
been  formed  in  past  geological  ages  from  the  bones  of  marine 
animals.  Phosphate  rock,  after  treatment  with  sulfuric  acid  as 
in  the  case  of  bone  ash,  is  used  in  enormous  quantities  as  a 
fertilizer. 

161.  Sodium  Carbonate  and  Bicarbonate. — The  carbonates 
of  sodium  which  were  referred  to  above  (156)  may  be  obtained 
by  passing  carbon  dioxide  into  sodium  hydroxide  solution;  we 
get  in  this  way  either  the  carbonate  Na2CO3,  or  the  bicarbonate 
NaHCO3,  according  to  the  proportion  of  carbon  dioxide  used. 
We  may  consider  that  the  gas  first  unites  with  water  to  form 
carbonic  acid,  which  then  reacts  with  sodium  hydroxide  according 
to  the  two  following  equations : 

2NaOH+H2CO3  ->  Na2CO3+  2H2O, 
NaOH+H2CO3  ->  NaHCO3+H2O. 

These  carbonates  of  sodium  are  manufactured  in  immense 
quantities,  as  they  are  very  important  substances.  In  practice 
they  are  not  made  according  to  the  reactions  given,  but  by  more 
economical  processes,  which  will  be  considered  later. 

162.  Potassium  Carbonate  and  Bicarbonate. — Potassium  also 
forms  analogous  carbonates,  K2GO3  and  KHCO3;   the  former, 
commonly  known  as  potash,  is  contained  in  wood  ashes,  from 
which  it  may  be  dissolved  by  water.    Upon  boiling  down  the 
solution  known  popularly  as  lye,  a  residue  of  crude  potassium 
carbonate,  K2C03,  remains.    This,  when  more  strongly  heated 
to  burn  out  brown  tarry  matters,  gives  white  potash,  so  called 
from  the  fact  that  the  evaporation  of  the  lye  is  carried  out  in  an 
iron  pot.    This  lye  is  extensively  used  in  the  preparation  of  a 
crude  soft  soap.    A  purer  form  of  potash  is  used  in  manufacturing 
liquid  soaps.    Common  hard  soap  is  made  from  sodium  car- 
bonate and  fats  of  various  kinds, 


Acids,  Bases,  and  Salts — //  97 

163.  The  Action  of  Acids  on  Carbonates. — If  some  hydro- 
chloric acid  is  poured  on  a  piece  of  marble  (150),  the  liquid 
appears  to"  boil,  although  the  temperature  does  not  rise  notice- 
ably.   It  is  easy  to  show  that  the  apparent  boiling,  called 
effervescence,  is  due  to  the  escape  of  carbon  dioxide  gas.    The 
marble  dissolves  completely  if  sufficient  acid  is  used,  and  the 
evaporated  solution  leaves  a  residue  of  calcium  chloride  (151). 
The  reaction  is  as  follows: 

CaCO3+2HCl->  CaCl2+H2O+CO2. 

Similar  reactions  take  place  between  calcium  carbonate  and 
nitric  and  sulfuric  acids: 

CaCO3+2HNO3->  Ca(NO3)2+H2O+CO2, 
CaCO3+H2S04-»  CaSO4+H20+CO2. 

In  fact,  the  carbonates  of  other  elements  all  show  this  kind  of  a 
reaction  with  these  acids;  for  example: 

NaHCO3+HCl->  NaCl+H2O+CO2, 
K2CO3+2HNO3-»  2KNO3+H2O+CO2. 

In  general,  carbonates  are  decomposed  by  acids. 

164.  Barium  Sulfate:    a  Test  for  Sulfates.— The  element 
barium  resembles  calcium  (150)  very  closely  in  its  behavior. 
Let  us  consider  just  one  of  its  reactions  at  present,  leaving  a 
study  of  the  others  until  a  later  time.    Barium  sulfate,  BaS04, 
is  a  white  solid  which  is  as  insoluble  in  water  as  glass;  barium 
chloride,  BaCl2,  is  about  as  soluble  as  common  salt.     If  we  pour 
some  sulfuric  acid  into  a  clear,  colorless  solution  of  barium 
chloride,  a  white  precipitate  of  barium  sulfate  forms  at  once: 

BaCl2+H2SO4->  BaSO4+2HCl. 

We  should  observe  the  similarity  of  this  equation  to  that  for  the 
action  of  sulfuric  acid  on  calcium  chloride  (153). 

If  we  add  to  a  solution  of  barium  chloride  a  solution  of  sodium 
sulfate,  or  of  magnesium  sulfate,  or  in  fact  of  any  sulfate  whatso- 
ever, a  precipitate  of  barium  sulfate  is  formed.  For  example, 
with  magnesium  sulfate  we  have 

BaCl2+MgSO4->  BaSO4+MgCl3. 


98  Introduction  to  General  Chemistry 

Unlike  other  common  barium  salts  the  sulfate  is  insoluble  not 
only  in  water  but  also  in  acids.  If  barium  chloride  is  added 
to  a  solution  and  a  white  precipitate  appears  which  has  these 
properties,  a  sulfate  is  proved  to  be  present. 

165.  Copper  and  Its  Compounds. — The  important,  familiar 
metal  copper  is  an  element.  We  have  already  learned  that  when 
heated  in  air  or  oxygen  it  unites  with  oxygen  to  form  copper 
oxide  (32,  33,  82),  a  black  solid: 

2Cu+O2->2CuO. 

Copper  oxide  reacts  with  the  corresponding  acids  to  form  the 
chloride,  nitrate,  and  sulfate,  thus: 

CuO+2HCl-»  CuCl2+H2O 
CuO+2HNO3->  Cu(NO3)2+H2O 
CuO+H2SO4-»  CuSO4+H2O. 

These  salts  are  all  easily  soluble  in  water,  giving  blue  solutions. 
It  will  be  recalled  that  calcium  oxide  unites  with  water  to 
form  the  hydroxide,  thus : 

CaO+H2O->Ca(OH)2. 

On  the  other  hand,  if  we  bring  copper  oxide  and  water  together, 
no  union  takes  place.  This  might  be  taken  to  indicate  that 
copper  hydroxide,  which  we  might  expect  to  have  the  formula 
Cu(OH)2,  cannot  be  formed  or  does  not  exist.  This,  however, 
is  not  the  case;  it  is  a  well-known  substance  which  is  easily 
obtained  in  another  way.  If  we  add  to  a  solution  of  copper  sul- 
fate a  solution  of  sodium  hydroxide,  a  blue  precipitate  of  copper 
hydroxide  forms.  This  is  a  blue  solid  which  is  very  nearly 
insoluble  in  water.  Its  formation  takes  place  thus : 

CuSO4+2NaOH->  Cu(OH)2+Na2SO4. 

We  may  also  get  copper  hydroxide  by  the  interaction  of  solutions 
of  copper  chloride  or  nitrate  with  sodium  hydroxide : 

CuCl2+2NaOH->  Cu(OH)a+2NaCl. 

If  the  copper  hydroxide  formed  in  the  last  reaction  is  heated 
by  boiling  the  mixture,  the  blue  precipitate  turns  black.     This 


Acids,  Bases,  and  Salts — //  99 

change  in  color  is  due  to  a  change  of  part  of  the  hydroxide  into 
the  oxide  and  water: 

Cu(OH)2->CuO+H2O. 

166.  The  Preparation  of  Difficultly  Soluble  Hydroxides.— 

Many  hydroxides  of  elements  are  nearly  insoluble  in  water.  In 
such  cases,  the  hydroxides  are  formed  from  a  solution  of  a  salt 
of  the  element  by  adding  sodium  hydroxide,  or  potassium 
hydroxide,  or  in  many  cases  ammonium  hydroxide,  as  illustrated 
by  the  following  equations : 

CaCl2+2KOH->  Ca(OH)a+2KCl, 
MgSO4+  2NaOH  ->  Mg(OH)2+Na2SO4. 

Such  difficultly  soluble  hydroxides  separate  as  precipitates,  which 
may  be  filtered  out. 

167.  Lead  and  Its  Compounds. — The  well-known  metal  lead 
is  an  element,  which  is  used  extensively  in  metallic  form  and  also 
in  the  form  of  compounds.    Lead  unites  with  oxygen  directly 
when  heated  in  air  or  oxygen.,  giving,  under  suitable  conditions, 
the  pale  yellow  oxide  PbO,  known  as  litharge.    This  oxide,  like 
those  of  magnesium,  zinc,  and  copper,  reacts  with  acids  to  form 
salts.    Thus  with  nitric  acid  we  get  lead  nitrate : 

PbO+  2HNO3  ->  Pb(NO3)2+H2O. 

This  salt  forms  large,  white  crystals  which  dissolve  readily  in 
water  to  form  a  colorless  solution.  If  hydrochloric  acid  is  added 
to  a  solution  of  lead  nitrate,  a  white  precipitate  of  lead  chloride 
is  obtained: 

Pb(NO3)2+2HCl->  PbCl2+2HNO3. 

Lead  chloride  is  only  slightly  soluble  in  cold  water,  but  is  much 
more  soluble  in  hot  water,  from  which,  upon  cooling,  it  separates 
again  in  white  needle-shaped  crystals.  Lead  chloride  is  also 
obtained  from  litharge  and  hydrochloric  acid : 

PbO-f-  2HC1  ->  PbCl2+H2O. 


ioo  Introduction  to  General  Chemistry 

Upon  adding  dilute  sulfuric  acid  to  a  solution  of  lead  nitrate, 
a  white  precipitate  of  lead  sulf ate  forms : 

Pb(NO3)2+H2SO4->  PbSO4+2HN03. 

Lead  sulf  ate  is  nearly  insoluble  in  hot  or  cold  water. 

Metallic  lead  is  acted  upon  very  slowly  by  hydrochloric  acid. 
With  the  cold  dilute  acid  no  appreciable  action  takes  place;  with 
boiling,  concentrated  acid,  a  very  slow  reaction  occurs,  thus : 

Pb-h2HCl->PbCl2+H2. 

By  methods  to  be  considered  later,  it  is  possible  to  prepare 
an  oxide  of  lead  containing  double  the  proportion  of  oxygen 
present  in  PbO,  namely  Pb02,  or  lead  dioxide.  This  oxide  does 
not  react  with  dilute  nitric  acid.  When  it  is  heated  with  hydro- 
chloric acid  it  gives  lead  chloride  and  chlorine : 

PbO2+4HCl->  PbCl2+2H2O+Cl2. 
If  we  compare  this  reaction  with  the  following, 
PbO+2HCl->  PbCl2+H20, 

we  see  that  the  excess  of  oxygen  in  Pb02  above  that  in  PbO  oxidizes 
the  hydrochloric  acid,  forming  water  and  setting  free  chlorine. 

Lead  acetate,  Pb(C2H3O2)2'3H2O,  is  formed  by  dissolving 
litharge,  PbO,  in  acetic  acid.  It  forms  colorless  prismatic  crys- 
tals, which  are  readily  soluble  in  water.  It  is  a  poisonous  salt 
and  is  called  sugar  of  lead  on  account  of  its  sweetish  taste. 

1 68.  Silver  and  Its  Compounds. — Silver  is  so  familiar  a 
metal  that  we  need  not  describe  its  properties.  It  is  an  element 
which  is  most  extensively  used  in  the  metallic  form,  but  which 
forms  several  compounds  of  great  practical  importance.  The 
metal  is  not  readily  acted  upon  by  dilute  acids,  with  the  excep- 
tion of  nitric  acid,  with  which  it  undergoes  a  complex  reaction 
represented  by  the  equation 

3Ag+4HN03->  3AgN03+NO+2H20. 

We  need  not  consider  this  reaction  critically  at  this  time, 
although  it  is  well  worth  careful  study;  but  note  that  silver 


Acids,  Bases,  and  Salts — //  101 

nitrate  is  an  easily  soluble  salt,  forming  a  colorless  solution.  The 
solid  salt  forms  large  white  crystals. 

169.  Silver  Chloride,  AgCl. — The  addition  of  hydrochloric 
acid  to  a  solution  of  silver  nitrate  produces  at  once  a  heavy  white 
precipitate  of  silver  chloride,  which  is  almost  insoluble  in  water: 

AgN03+HCl->  AgCl+HN03. 

By  adding  an  excess  of  hydrochloric  acid  practically  all  of  the 
silver  in  a  solution  is  precipitated.  The  precipitate  does  not 
dissolve  appreciably  in  any  of  the  common  acids.  It  is,  however, 
very  easily  soluble  in  ammonia  solution,  from  which  it  is  again 
thrown  down  if  the  solution  is  acidified  with  nitric  or  hydrochloric 
acid.  If  any  solution  of  unknown  nature  gives  with  hydro- 
chloric acid  a  white  precipitate  which  is  insoluble  in  an  excess  of 
the  acid,  but  easily  soluble  in  ammonia,  from  which  solution  it  is 
thrown  down  by  acidifying  the  solution  with  hydrochloric  acid, 
it  is  safe  to  conclude  that  the  original  solution  contained  a  salt 
of  silver.  This  series  of  reactions  constitutes  a  test  for  silver  in 
the  form  of  a  dissolved  salt. 

170.  Silver  Sulfate,  Ag2SO4.— This  salt  is  formed  as  a  white 
crystalline  precipitate  when  sulfuric  acid  is  added  to  a  con- 
centrated solution  of  silver  nitrate.     It  is  not  very  soluble,  i  g. 
requiring  about  2oac.c.  of  cold  water  for  its  solution.     The  same 
salt  is  also  formed  by  the  action  of  hot,  concentrated  sulfuric 
acid  on  metallic  silver. 

171.  Silver  Phosphate,  Ag3PO4. — This  salt  is  formed  as  a 
yellow  precipitate  when  sodium  phosphate,  Na3PO4,  or  some 
other  soluble  phosphate  is  added  to  a  solution  of  silver  nitrate: 

3  AgN03+Na3P04  ->  Ag3P04+3NaNO3. 

This  yellow  precipitate  is  readily  soluble  in  dilute  nitric  acid, 
forming  a  colorless  solution.  It  also  dissolves  easily  in  aqueous 
ammonia,  giving  a  colorless  solution,  from  which  it  is  again 
thrown  down  when  the  solution  is  exactly  neutralized  with  nitric 
acid. 

Silver  may  be  distinguished  from  lead  most  easily  by  reason 
of  the  solubility  of  lead  chloride  in  hot  water,  in  which  silver 
chloride  is  insoluble. 


102  Introduction  to  General  Chemistry 

172.  Silver  Oxide,  Ag2O. — The  addition  of  sodium  hydroxide 
to  a  solution  of  silver  nitrate  gives  a  black  precipitate  of  silver 
oxide,  Ag2O.    We  might  expect  silver  hydroxide,  AgOH.  to  be 

formed  thus: 

AgNO3+NaOH-»  AgOH+NaNO3. 

Possibly  this  is  what  first  happens,  but,  if  so,  the  hydroxide 
formed  changes  at  once  into  the  oxide, 

2AgOH->Ag20+H2O. 

It  will  be  recalled  that  copper  hydroxide  is  decomposed  at  the 
temperature  of  boiling  water  into  the  oxide  and  water  (165). 
In  the  case  of  silver  hydroxide  the  change  takes  place  at  room 
temperature. 

173.  Iron  and  Its  Compounds. — The  element  iron  is  the  most 
important  of  all  metals.    It  unites  directly  with  oxygen  at  a  red 
heat,  forming  the  oxide  Fe304  (81).    It  can  also  form  two  other 
oxides,  FeO  and  Fe203.    The  oxide  Fe3O4  is  magnetic  and  is 
called  magnetic  iron  oxide;   FeO  is  called  ferrous  oxide  (from 
ferrum,  iron),  while  Fe2O3  is  called  ferric  oxide.    Ferrous  oxide 
gives,  with  the  corresponding  acids,  ferrous  chloride,  FeCl2,  and 
ferrous  sulfate,  FeS04.    These  salts  are  also  formed  from  iron 
by  the  following  reactions:  „ 

Fe+2HCl->FeCl2+H2, 
Fe+H2SO4-»FeSO4+H2. 

In  all  ferrous  compounds  the  valence  of  iron  is  two. 

The  action  of  hydrochloric  acid  on  ferric  oxide  takes  place 
thus: 

Fe2O3+6HCl-»  2FeCl3+3H2O. 

The  salt  FeCl3  is  called  ferric  chloride.  It  is  a  dark-yellow 
substance  which  dissolves  easily  in  water  to  form  a  yellow  solu- 
tion. On  the  other  hand,  ferrous  chloride,  FeCl2,  is  pale  green 
and  forms  a  pale-green  solution.  We  cannot  get  FeCl3  from 
iron  and  hydrochloric  acid,  but  we  do  get  the  salt  by  the  action 
of  chlorine  on  ferrous  chloride, 

2FeCl2+Cl3->2FeCl3, 


Acids,  Bases,  and  Salts — II  103 

or  on  iron, 

2Fe+3Cl2->2FeCl3. 

In  ferric  chloride  the  valence  of  the  iron  is  three.  We  also  know 
ferric  nitrate,  Fe(N03)3,  and  ferric  sulfate,  Fe2(S04)3.  There  are, 
therefore,  two  series  of  iron  salts — the  ferrous,  in  which  the  valence 
of  iron  is  two,  and  the  ferric,  in  which  the  "valence  is  three. 

We  can  obtain  Ihe  two  hydroxides  of  iron,  both  of  which  are 
nearly  insoluble  in  water,  by  the  action  of  sodium  hydroxide  on 
solutions  of  ferrous  and  ferric  salts: 

FeCl2+ 2NaOH  ->  Fe(OH)2+  2NaCl, 
FeCl3+3NaOH-»  Fe(OH)3+3NaCl. 

Ferrous  hydroxide  is  white  if  pure,  but  is  usually  obtained  as  a 
dirty-green  precipitate;  this  is  due  to  partial  oxidation  by  the 
-action  of  oxygen  of  the  air,  with  which  it  readily  unites.  Ferric 
hydroxide  is  a  brown  precipitate.  These  hydroxides  unite  with 
acids  to  form  salts: 

Fe(OH)2+  2HC1  ->  FeCl2+  2H2O, 
Fe(OH)3+3HCl->  FeCl3+3H20. 

174.  Aluminum  and  Its  Compounds. — The  common  metal 
aluminum  is  an  element.  As  is  well  known,  the  metal  is  not 
acted  upon  by  air  or  water.  It  reacts  easily  with  dilute  hydro- 
chloric acid,  giving  aluminum  chloride,  A1C13,  and  hydrogen: 

2Al+6HCl->  2A1C13+3H2. 

Upon  evaporation,  the  solution  deposits  white  crystals  of  the 
compound  A1C13'6H2O.  It  is  not  possible  to  obtain  the  anhy- 
drous salt,  A1C13  by  heating  these  crystals,  for  the  purpose  of 
driving  off  water,  since  they  decompose  thus : 

2A1C1S-  6H2O  ->  A12O3+6HC1+3H2O. 

The  anhydrous  chloride  is  formed  by  the  action  of  dry  chlorine 
gas  on  aluminum : 

2Al+3Cl«->2AlClj. 


104  Introduction  to  General  Chemistry  . 

Aluminum  chloride  is  easily  soluble  in  water,  forming  a  colorless 
solution.  This  solution  gives  with  ammonia  a  white  precipitate 
of  aluminum  hydroxide,  A1(OH)3,  which  is  insoluble  in  water: 

A1C13+3NH4OH  -»  A1(OH)3+3NH4C1. 
When  heated,  the  hydroxide  gives  the  oxide  and  water: 
2  A1(OH) 


Rubies  and  sapphires  are  natural  forms  of  aluminum  oxide. 
Emery,  which  is  a  valuable  abrasive,  is  an  impure  form  of  the 
same  substance. 

175.  Various  Aluminum  Salts.  —  The  hydroxide  is  a  base 
which  reacts  with  acids  to  give  the  corresponding  salts,  thus: 

Al(OH)3+3HCl->  A1C13+3H20, 
A1(OH)3+3HN03  ->  A1(N03)3+3H20, 
2A1(OH)3+3H2S04  ->  A12(S04)3+6H20. 

The  nitrate  and  sulfate  are  easily  soluble  in  water,  giving  color- 
less solutions.  The  well-known  substance  alum  is  potassium, 
aluminum  sulfate,  KAl(SO4)2-i2H2O.  It  is  obtained  in  large, 
colorless  crystals  when  a  solution  made  from  potassium  sulfate 
and  aluminum  sulfate  is  allowed  to  evaporate.  The  correspond- 
ing sodium  and  ammonium  salts  are  well  known,  and  have  the 
formulae  NaAl(SO4)2-i2H2O  and  NH4Al(S04)2-i2H2O,  respec- 
tively. All  such  compounds  are  known  as  double  salts  ;  chemists 
are  familiar  with  a  great  variety  of  these.  Other  examples  of 
well-known  double  salts  are  ammonium  ferrous  sulfate,  (NH4)2 
Fe(SO4)2-6H20,  and  potassium  cupric  chloride,  K2CuCl4'2H20. 

176.  Acid  Reaction  of  Aluminum  Salts.  —  Solutions  of  the 
chloride,  nitrate,  and  sulfate  of  aluminum,  and  also  of  alum,  are 
not  neutral,  as  we  might  expect,  but  are  distirctly  acid  in  reaction. 
They  also  have  a  sour  taste.    On  the  other  hand,  we  find  that 
moist  aluminum  hydroxide,  if  it  has  been  carefully  washed  free 
from  the  ammonia  used  in  precipitating  it,  has  no  action  on  either 
blue  or  red  litmus.    It  is  also  tasteless.    Nevertheless,  we  call 
the  hydroxide  a  base,  because  it  unites  with  acids  to  form  salts. 
We  say,  however,  that  it  is  a  weak  base  ;  and  we  find  in  general 


Acids,  Bases,  and  Salts — II  105 

that  weak  bases,  of  which  many  are  known,  give  salts  whose 
solutions  are  acid  in  reaction.  This  is  an  important  matter  which 
will  have  to  be  studied  carefully  later. 

177.  Acid  Properties  of  Aluminum  Hydroxide. — If  we  add 
sodium  hydroxide  to  a  solution  of  an  aluminum  salt,  a  white 
precipitate  of  aluminum  hydroxide  is  first  formed,  just  as  with 

ammonia : 

AlCl3+3NaOH->Al(OH)3+3NaCl. 

However,  upon  adding  an  excess  of  sodium  hydroxide,  we  find 
that  the  precipitate  goes  into  solution.  If  pure  aluminum  hy- 
droxide is  dissolved  in  a  solution  of  sodium  hydroxide  and  the 
resulting  solution  evaporated,  crystals  of  sodium  aluminate, 
NaAlO2,  are  obtained.  This  substance  is  easily  soluble  in  water 
and  is,  in  reality,  a  salt.  It  thus  appears  that  aluminum  hy- 
droxide acts  as  an  acid  in  this  case,  and  we  might  write  the  equa- 
tion for  the  action  of  sodium  hydroxide  upon  it  thus: 

HA102-H20+NaOH->  NaA102+2H2O. 

We  find  that  the  solution  of  sodium  aluminate  is  strongly  alkaline 
toward  litmus,  and  say,  therefore,  that,  although  aluminum 
hydroxide  has  some  acid  properties,  it  is  a  very  weak  acid. 

Thus  we  see  that  a  substance  may  be  both  a  base  and  an  acid. 
Such  a  substance  is  said  to  be  amphoteric.  Several  metallic 
hydroxides  are  amphoteric.  Thus  zinc  hydroxide,  Zn(OH)2, 
forms  with  hydrochloric  acid,  ZnCl2,  and  with  sodium  hydroxide, 
Na2Zn02,  sodium  zincate.  It  is  of  interest  to  note  that  aluminum 
hydroxide  does  not  react  with  carbonic  acid,  and  in  fact  no  car- 
bonate of  aluminum  has  ever  been  made.  Now,  carbonic  acid 
is  a  very  weak  acid,  and  aluminum  hydroxide  is  a  very  weak  base. 
In  general,  we  find  that  -very  weak  bases  do  not  form  salts  with 
very  weak  acids. 

Compounds  of  aluminum  are  very  abundant  in  the  earth. 
Common  clay  and  numerous  kinds  of  common  rocks  are  com- 
pounds of  aluminum. 

178.  Mercury  and  Its  Compounds. — We  have  already  learned 
something  of  the  chemical  behavior  of  mercury  and  mercuric 
oxide,  HgO  (13, 14,  86).    The  oxide,  which  is  insoluble  in  water, 


io6  Introduction  to  General  Chemistry 

dissolves  in  dilute  hydrochloric  acid,  giving  mercuric  chloride, 
HgCl2,  and  in  nitric  acid,  giving  mercuric  nitrate,  Hg(NO3)2: 

HgO+2HCl->  HgCl2+H20. 
HgO+  2HNO3  ->  Hg(N03)2+H20. 

These  salts  form  white  crystals  which  are  soluble  in  water.  The 
soluble  salts  of  mercury  are  all  extremely  poisonous  when  taken 
internally.  Mercuric  chloride  is  familiarly  j|own  as  bichloride 
of  mercury  or  corrosive  sublimate,  and  is  extensively  used  as  a 
powerful  germicide  and  antiseptic. 

179.  The  Formation  of  Mercuric  Salts. — The  nitrate  can  be 
made  by  the  action  of  warm,  concentrated  nitric  acid  upon 
metallic  mercury: 

3Hg+8HN03->  3Hg(N03)2+2NO+4H20. 

Hydrochloric  acid  does  not  act  appreciably  upon  mercury;  but 
the  chloride  can  be  obtained  by  the  action  of  chlorine  on  the 

metal : 

Hg+Cl2->HgCl2. 

It  is  also  made  by  heating  a  mixture  of  mercuric  sulfate  and 
common  salt: 

HgSO4+2NaCl  ^HgCl2+Na2SO4. 

The  mercuric  chloride  formed  is  readily  volatile  and  is  separated 
by  sublimation;  hence  the  old  name  "corrosive  sublimate." 
The  process  of  vaporization  of  a  solid  and  the  condensation  of  its 
vapor  directly  to  the  crystalline  form  is  called  sublimation. 
The  sulfate  is  made  by  strongly  heating  mercury  with  concen- 
trated-sulf  uric  acid: 

Hg+  2H2SO4  ->  HgSO4+  SO2+  2H2O. 

1 80.  Mercurous  Salts. — The  action  of  cold,  dilute  nitric  acid 
on  an  excess  of  mercury  gives  rise  to  a  solution  of  a  salt  having  the 
formula  HgNO3  and  called  mercurous  nitrate.    A  solution  of 
this  salt  gives  with  hydrochloric  acid  a  white  precipitate  of 
mercurous  chloride,  HgCl,  and  with  dilute  sulf  uric  acid  also  a 
white  precipitate  of  mercurous  sulfate.  Hg2S04;   both  of  these 


Acids,  Bases,  and  Salts — //  107 

precipitates  are  practically  insoluble  in  water.  Thus  mercury, 
like  iron,  forms  two  series  of  salts:  tne  mercurous,  in  which  the 
element  has  a  valence  of  one,  or  is  univalent,  and  the  mercuric, 
in  which  it  has  a  valence  of  two,  or  is  bivalent. 

181.  The  Two  Oxides  of  Mercury. — A  solution  of  mercuric 
nitrate  gives  with  a  solution  of  sodium  hydroxide  a  yellow 
precipitate  of  mercuric  oxide,  HgO: 

Hg(NO3)2V  2NaOH  ->  HgO+  2NaNO3+H2O. 

The  hydroxide  of  mercury,  like  that  of  silver,  cannot  be  obtained; 
we  might  say  that  it  is  so  unstable  that  it  changes  into  the  oxide 
and  water  as  soon  as  it  is  formed;  in  this  respect  it  resembles 
the  corresponding  compound  of  silver.  The  yellow  oxide, 
formed  in  this  way,  seems  to  differ  from  the  red  oxide,  obtained 
by  heating  mercury  in  the  air  or  in  oxygen,  only  in  being  made 
up  of  very  much  smaller  particles. 

A  solution  of  mercurous  nitrate  gives  with  sodium  hydroxide 
a  nearly  black  precipitate  of  mercurous  oxide,  Hg2O. 

182.  Calomel. — Mercurous  chloride,  HgCl,  which  is  com- 
monly called  calomel,  is  extensively  used  in  medicine.     It  is  a 
remarkable  but  well-known  fact  that  the  usual  Adi^hal  dose 
of  calomel  contains  many  times  as  much  mercury  as  does  a  fatal 
dose  of  mercuric  chloride.    This  great  difference  in  physio- 
logical effect  is  in  part  due  to  the  fact  that  while  mercuric 
chloride  is  easily  soluble  in  water,  mercurous  chloride  is  nearly 
insoluble.  $. 

183.  The  Valencies  of  Radicals. — Now  that  we  have  studied 
a  considerable  additional  number  of  acids,  bases,  and  salts, 
we  may  again  revert  to  the  study  of  valence  (146, 147)  since  it 
furnishes  a  key  to  the  easy  mastery  of  formulae,  an  undertak- 
ing which  is  as  necessary  to  the  study  of  chemistry  as  learning  to 
spell  is  in  the  mastery  of  a  language.     To  write  the  formula  of  a 
chloride  of  a  metal  it  is  only  necessary  to  group  with  the  symbol 
of  the  latter  as  many  chlorine  symbols  as  the  metal  in  question 
has  units  of  valence;  thus  the  formulae  of  barium  chloride  and 
of  aluminum  chloride  are  BaCl2  and  A1C13  respectively.     If  we 
wish  to  write  the  formulae  of  the  nitrates  we  group  the  nitrate 


io8 


Introduction  to  General  Chemistry 


radical  with  the  metal  symbol  in  question  according  to  the  same 
rule,  thus  Ba(NO3)2  and  A1(NO3)3.  To  write  the  formulae  of 
sulfates,  we  must  again  group  the  symbols  of  the  radical  and 
the  metal  so  that  the  total  valence  of  each  satisfies  that  of  the 
other.  Since  the  sulfate  radical  is  bivalent,  we  may  have  to 
use  more  than  one  symbol  weight  of  either  the  sulfate  or  the 
metal.  Thus  the  formula  of  sodium  sulfate  is  Na2S04;  that  of 
barium  sulfate  is  BaSO4;  while  that  of  aluminum  sulfate  is 
Al^SO^  Since  PO4  is  trivalent,  we  know  at  once  that  the 
Tormulaof  sodium  phosphate  must  be  Na3PO4,  that  of  barium 
phosphate  must  be  Ba3(PO4)2,  and  that  of  aluminum  A1PO4. 
These  examples  are  sufficient  to  show  how  a  knowledge  of 
valence  simplifies  the  writing  of  formulae.  In  Table  IX  the 

TABLE  IX 


H.  . 

Cl 

HC1 

HOH  (H2O) 

Na 

NO, 

NaCl 

HNO~ 

I 

K.    . 

OH 

HC2H3O2 

NaOH 

Univalent 

NIL.  . 

C2H3Oa 

NH4C1 

NH4OH 

Ag.  . 

AgCl 

Ag2O 

Hg.. 

HgCl 

Na2SO4 

f  * 

Mg  
Ca  

S04 
CO, 

MgCl2 
CaCl2 

H2S04 
H2CO3 

Ba 

o 

BaCl2 

CaCO3 

II 

Zn 

ZnCl2 

BaSO4 

Bivalent 

Fe..  /..... 
Cu 

FeCl2 
CuCl2 

CuS04 
Ca(OH)a 

Pb 

PbCl2 

MgO 

Hg 

HgCl2 

HgO 

III 

Al 

PO4 

A1C13 

H,PO4 

Trivalent 

Fe./.::::: 

N 

FeCl 

A1P04 
NH3 

IV 

c 

CC14 

CH4 

Quadrivalent 

CO2 

various  elements  and  radicals  studied  are  classified  with  respect 
to  their  valencies,  which  vary  from  one  to  four.  In  the  column 
headed  by  hydrogen  we  have  the  metals  together  with  the 
ammonium  radical.  In  the  next  column,  headed  by  chlorine, 
are  the  elements  and  radicals  that  unite  as  a  rule  with 
those  of  the  first  column.  The  last  two  columns  contain  the 
formulae  of  some  typical  compounds. 


CHAPTER  X 

THE  KINETIC  THEORY  OF  MATTER  AND  THE  MOLECULAR 
HYPOTHESIS 

184.  An  Old  Greek  Hypothesis. — The  question  whether  a 
portion  of  a  given  substance,  say  a  drop  of  water,  could  be  sub- 
divided to  an  unlimited  extent,  and  whether  the  smallest  particle 
so  produced  would  still  differ  in  no  way  except  in  size  from  the 
original,  was  one  which  was  much  debated  by  the  Greek  philos- 
ophers centuries  before  chemistry  became  a  science.    Anaxagoras 
(B.C.  500),  who  held  that  there  was  no  limit  to  the  divisibility 
of  matter,  was  opposed  by  Democritus  (B.C.  470),  who  taught 
that  in  the  imagined  process  of  continued  subdivision  minute 
particles  would  finally  be  encountered  which  could  not  be  cut 
in  two  without  destroying  or  completely  changing  the  nature 
of  the  substance;   these  particles  were  called  atoms  (arojuos,  a 
body  which  cannot  be  cut  in  two).    This  idea  may  be  illustrated 
by  the  following  analogy.    If  we  take  a  bushel  of  wheat,  we  may 
divide  it  into  pecks,  quarts,  gills,  etc.,  and  yet  each  measure  of 
the  material  will  be  a  quantity  of  wheat;  we  may  go  still  farther, 
but  we  will  ultimately  reach  the  single  grains,  which  are  still 
grains  of  wheat;  but  if  we  cut  these  grains  in  two  the  resulting 
parts  may  no  longer  be  called  wheat,  since  they  would  no  longer 
possess  the  most  remarkable  property  of  wheat,  which  is  that  of 
growing  if  planted.    At  present  we  use  the  term  molecule  (from 
Latin  molecula,  the  diminutive  of  moles,  a  mass)  to  mean  essen- 
tially the  same  as  the  term  "atom,"  as  used  by  the  Greeks,  and 
speak  therefore  of  the  molecular  theory  of  matter  and  the  molec- 
ular hypothesis.    The  present  chapter  gives  an  account  of 
this  hypothesis  and  aims  to  show  how  it  furnishes  us  an  explana- 
tion of  many  important  facts. 

185.  The  Molecules  of  Water.— According  to  the  molecular 
hypothesis,  a  drop  of  water  can  be  subdivided,  and  still  remain 
water,  only  until  the  single  molecules  are  reached,  and  no  further. 
The  splitting  up  of  the  molecules  of  water  might  separate  the 

109 


no  Introduction  to  General  Chemistry 

smallest  particles  of  oxygen  from  those  of  hydrogen;  but  the 
result  would  be  the  decomposition  of  the  water  into  its  elementary 
constituents.  The  particles  of  the  elements  oxygen  and  hydro- 
gen of  which  the  molecules  of  water  are  made  up  are  now  called 
atoms.  It  is  supposed  that  all  of  the  molecules  of  water  are 
alike  in  every  respect,  each  being  made  up  of  one  atom  of  oxygen 
and  two  atoms  of  hydrogen.  The  reason  for  this  last  conclusion 
is  discussed  in  the  next  chapter.  In  general,  each  molecule  of  a 
given  pure  substance  is  just  like  every  other  molecule  of  that  sub- 
stance. The  nature  of  a  substance  is  determined  by  the  nature 
of  its  molecules;  and  this  is,  in  turn,  determined  by  the  number 
and  kind  of  atoms  composing  the  molecules.  We  imagine  mole- 
cules to  be  very  small,  since  they  cannot  be  seen  with  the  aid 
of  a  microscope  of  the  highest  power.  A  cubic  centimeter  of  air 
may  contain  an  almost  inconceivable  number  of  these  tiny 
particles. 

186.  The  Molecular  Hypothesis  Applied  to  Gases. — Let  us 
first  consider  the  known  facts  concerning  gases  and  try  to  see  how 
these  facts  can  be  connected  with  the  supposition  that  gases  are 
made  up  of  small  particles,  the  molecules.    In  the  first  place, 
we  know  that  a  confined  gas  tends  to  expand  and  exerts  a  pressure 
on  the  walls  of  the  vessel  which  holds  it.     If  we  increase  the 
pressure,  the  volume  of  the  gas  is  diminished,  and  by  applying 
a  great  pressure  the  decrease  in  volume  may  be  made  very  great. 
We  may  explain  this  in  either  of  two  ways :  first,  that  the  mole- 
cules, like  rubber  balloons,  are  themselves  compressible;    or, 
second,  that  the  molecules,  which  may  not  be  appreciably  com- 
pressible, are  at  considerable  distances  from  one  another,  but 
are  brought  closer  together  when  the  gas  as  a  whole  is  compressed. 
Let  us  follow  up  this  second  idea  and  see  to  what  it  leads.    Two 
important  questions  now  present  themselves:   (i)  Why  are  the 
molecules  not  in  contact — that  is,  why  should  they  be  at  con- 
siderable distances  from  one  another?     (2)  Why  does  a  gas  exert 
a  pressure  in  all  directions  on  the  walls  of  the  vessel  which  con- 
tains it? 

187.  Are  Molecules  at  Rest  or  in  Motion?: — Would  it  make  a 
difference  in  the  state  of  affairs  whether  the  molecules  were  at 


Molecular  Hypothesis  in 

rest  or  in  motion?  Suppose  they  are  in  rapid  motion:  what 
would  follow?  Let  us  recall  Newton's  first  law  of  motion: 
A  body  at  rest  remains  at  rest,  and  a  body  in  motion  continues  to 
move  with  constant  velocity  in  a  straight  line,  unless  acted  upon 
by  some  external  unbalanced  force.  Now,  if  molecules  are  in 
motion,  and  if  they  behave  in  the  same  manner  as  other  bodies, 
and  if,  further,  they  are  elastic — that  is,  if  they,  like  rubber  or 
ivory  balls,  can  rebound  from  one  another  or  from  the  walls  of 
the  containing  vessel,  then  they  will  tend  to  continue  in  motion. 
Of  course,  in  such  a  case  as  the  one  imagined,  the  molecules  of  the 
gas  would  very  frequently  strike  one  another  and  also  the  walls  of 
the  containing  vessel;  but,  being  elastic,  they  would  rebound  and 
continue  in  motion  in  a  new  direction;  and  although  the  velocity 
of  an  individual  molecule  might  be  increased  or  decreased  as  the 
result  of  a  collision  with  another  molecule,  on  the  whole  the 
average  velocity  of  all  the  molecules  would  be  constant. 

The  various  elaborations  of  the  ideas  here  presented 
constitute  the  Kinetic  Theory  of  Matter.  This  hypothesis  is 
the  most  important  corollary  of  the  molecular  hypothesis.  It 
has  proved  enormously  fruitful  in  explaining  very  diverse 
phenomena  and  in  suggesting  new  lines  of  investigation. 

188.  The  Cause  of  Gas  Pressure;  Boyle's  Law.— The  strik- 
ing of  a  gas  molecule  against  the  wall  of  the  vessel  would  deliver  a 
little  blow;  and  if  millions  of  molecules  struck  each  square  centi- 
meter every  second  the  effect  would  be  to  tend  to  push  back  the 
surface.  But  what  is  this  but  the  exertion  of  pressure  ?  If  the 
molecules  strike  often  enough,  regularly  enough,  and  close  enough 
together,  this  pressure  would  seem  constant  and  uniform.  Now, 
suppose  the  gas  to  be  compressed  until  it  occupies  half  its  original 
volume.  In  each  cubic  centimeter  there  would  now  be  double 
the  original  number  of  molecules,  and  on  each  square  centimeter 
of  the  wall  of  the  container  twice  as  many  molecules  would 
strike  per  second  as  before;  so  that,  as  the  mass  of  each  mole- 
cule and  also  its  velocity  have  remained  unchanged,  we  should 
expect  just  double  the  pressure  per  square  centimeter;  and,  in 
fact,  this  is  just  what  we  find  by  experiment.  Thus  we  see  that 
by  imagining  a  gas  to  be  made  of  numerous  small,  rapidly  moving, 


H2  Introduction  to  General  Chemistry 

elastic  particles,  the  molecules,  we  get  an  explanation  of  gas 
pressure  and  of  Boyle's  law.  We  also  see  how  it  would  be  pos- 
sible for  the  molecules  to  be  at  considerable  distances  (com- 
pared with  their  own  diameters)  from  one  another  without 
tending  to  fall  together  into  a  mass  in  which  the  molecules  would 
all  be  permanently  in  contact. 

189.  The  Effect  of  Temperature  on  Molecular  Velocity. — We 
may  next  consider  why  it  is  that  the  molecules  are  in  motion  and 
whether  the  average  velocity  of  the  molecules  of  a  given  gas 
can  ever  be  changed.    We  know,  of  course,  that,  for  a  constant 
volume,  the  pressure  exerted  by  a  gas  increases  with  rise  of 
temperature,  so  that  if  we  are  to  explain  gas  pressure  as  due  to  the 
momenta  (mass  X  velocity)  of  the  molecules  which  strike  the  walls 
of  the  container,  we  must  suppose  that  an  increase  of  tempera- 
ture increases  either  the  mass  of  a  molecule  or  its  velocity,  or 
both.    Naturally,  the  number  of  collisions  with  the  wall  could 
not  be  increased  unless  the  velocity  increased.    Now,  it  would 
seem  more  reasonable  to  think  of  rise  of  temperature  as  causing 
an  increase  in  velocity  than  to  think  of  it  as  causing  an  increase 
in  mass;  so  we  have  only  to  imagine  that  a  rise  in  temperature 
causes  the  average  "velocity  of  the  molecules  to  increase  in  order  to 
get  a  simple  and  satisfying  explanation  of  the  effect  of  tempera- 
ture on  the  pressure  of  a  gas. 

On  the  other  hand,  a  decrease  in  temperature  is  accom- 
panied by  a  decrease  in  pressure;  and,  indeed,  the  pressure  is  at 
all  times  proportional  to  the  absolute  temperature.  This  would 
imply  that  the  pressure  would  be  zero  at  the  absolute  zero  of 
temperature.  But  zero  pressure  could  only  result  if  the  mole- 
cules were  completely  at  rest.  We  may  suppose,  therefore,  that 
at  absolute  zero  there  is  no  molecular  motion.  A  body  when  hot 
differs  from  the  same  body  when  cold  only  by  reason  of  the  more 
rapid  motion  of  its  molecules.  In  short,  according  to  this  way 
of  looking  at  the  matter,  heat  is  merely  the  outward  manifestation 
of  molecular  motion. 

190.  The    Mixing    of    Gases. — We  have   already    learned 
(122-124)  that,  as  a  rule,  liquids  and  solids  do  not  form  perfect 
mixtures  (solutions)  in  all  proportions;  that  is  to  say,  a  solid  or  a 


Molecular  Hypothesis  JI3 

liquid  will  dissolve  in  a  second  liquid  only  to  a  limited  extent. 
Not  so  with  gases :  every  gas  will  form  with  any  proportion  of 
any  other  gas  a  perfectly  uniform  mixture.  Air,  for  example,  is 
a  perfectly  homogeneous  mixture  of  several  gases.  Of  course,  if 
some  of  the  gases  which  are  brought  together  react  chemically, 
new  liquid  or  solid  compounds  might  be  formed  which  would 
separate  from  the  gaseous  mixture.  But  for  gases  that  do 
not  react  chemically  we  find  that  all  gases  mix  perfectly  in  all 
proportions. 

191.  The  Diffusion  of  Gases. — If  we  bring  two  gases  into,  the 
same  vessel  without  attempting  to  mix  them  we  find,  after  a  time, 
that  a  perfectly  uniform  mixture  is  present  in  the  vessel.     We  also 
know  that  if  a  gas  like  ammonia  is  liberated  at  one  place  in  a 
closed  room  its  odor  is  soon  perceptible  everywhere  in  the  room. 
The  process  of  the  spontaneous  mixing  of  gases  is  called  diffusion, 
and  we  say  that  the  ammonia  has  diffused  through  the  air  of  the 
whole  room.     It  is  now  easy  to  understand  how  this  diffusion 
takes  place.     The  molecules  of  ammonia  are  moving  in  all 
directions  with  high  velocities;  the  same  is  true  also  of  the  mole- 
cules of  the  air;    their  complete  and  uniform  intermingling  is 
therefore  inevitable. 

192.  The  Law  of  Partial  Gas  Pressures. — It  is  easily  found  by 
experiment  that,  if  portions  of  various  gases  are  brought  together 
in  the  same  vessel,  the  total  pressure  exerted  by  the  gas  mixture  is 
the  sum  of  the  pressures  that  would  be  exerted,  at  the  same  tempera- 
ture, by  the  same  portions  of  these  gases  if  each  occupied  the  space 
alone.    This  is  known  as  Dalton's  Law  of  Partial  Pressures. 
It  is  not  difficult  to  explain  this  law,  since,  in  a  mixture  of  gases, 
the  molecules  of  a  given  sort  will  strike  a  given  area  of  the  wall 
just  as  often  in  the  presence  of  other  unlike  molecules  as  in  their 
absence.     Each  kind  of  molecule  will  therefore  produce  the  same 
partial  pressure  as  if  the  others  were  absent: 

193.  Avogadro's  Hypothesis. — The   student  must   already 
have  been  impressed  by  the  fact  that  all  gases  show  great  similarity 
in  physical  behavior.    They  all  conform  to  the  laws  of  Boyle  and 
Charles.    This  fact,  together  with  others  which  we  shall  consider 
later,  led  Avogadro,  then  professor  of  physics  in  Turin,  Italy,  to 


ii4  '  Introduction  to  General  Chemistry 

suggest  in  1811  that  it  is  probable  that  all  gases  contain  the  same 
number  of  molecules  per  cubic  centimeter.  Although  this  sugges- 
tion received  some  support  for  the  first  twenty  years  after  its 
proposal,  it  was  then  nearly  forgotten  until  about  1860,  since 
which  time  its  importance  and  probability  have  been  impressed 
more  and  more  deeply  on  the  minds  of  physicists  and  chemists, 
so  that  during  the  last  fifty  years  it  has  become  one  of  the  most 
fundamental  principles  of  chemistry.  It  may  be  stated  concisely 
thus :  Equal  volumes  of  every  gas  or  vapor  at  the  same  temperature 
and  pressure  contain  the  same  number  of  molecules. 

Further  reasons  for  accepting  Avogadro's  hypothesis  will 
be  given  in  the  next  chapter.  Indeed,  the  evidence  from  so 
many  independent  sources  for  the  truth  of  this  view  is  now  so 
convincing  that  the  hypothesis  is  looked  upon  by  many  as  a 
statement  of  fact,  and  in  consequence  is  referred  to  as  Avogadro's 
Law. 

194.  Gas  Statistics. — Within  the  last  few  years  methods  have 
been  found  by  means  of  which  the  number  of  molecules  in  i  c.c. 
of  a  gas  has  been  found  with  a  high  degree  of  probability  and 
accuracy.    At  o°  and  76  cm.,  i  c.c.  of  any  gas  has  been  found  to 
contain  2.7X  io19  molecules,  with  a  probable  error  of  less  than  i 
per  cent.    The  number  of  molecules  in  22 . 4  liters  of  a  gas  under 
standard  conditions  is  therefore  22,4ooX2.7XioI9  =  6.o6Xio23. 
The  number  of  molecules  in  i  c.c.,  twenty-seven  millions  of  mil- 
lions of  millions,  is  so  immense  that  it  is  difficult  for  the  mind  to 
get  any  tangible  conception  of  its  magnitude.     However,  if  we 
think  of  the  molecules  in  i  c.c.  as  at  rest  for  the  moment,  and 
uniformly  distributed  in  rows  and  layers,  we  should  then  have 
in  each  row  of  i  cm.  length  ^27  X  iol8  =  3  X  io6,  or  three  million 
molecules,  a  number  which,  although  large,  is  at  least  compre- 
hensible.   Then  in  each  layer  there  would  be  three  million  of 
these  rows,  and,  in  the  whole  cubic  centimeter,  three  million  such 
layers. 

195.  A  Cubic  Mile  of  Sand. — Another  mental  picture  of  the 
case  may  be  got  if  we  imagine  i  c.c.  of  gas  to  have  been  expanded 
until  it  occupied  a  cubic  mile.    Then  each  row  of  molecules 
would  be  a  mile  long  and  would  contain  three  million  molecules, 


Molecular  Hypothesis  115 

spaced  about  TV  of  an  inch  apart.  Now,  a  grain  of  fine  sand  is 
about  yV  of  an  inch  in  diameter;  and  three  million  such  grains 
placed  side  by  side  would  extend  one  mile.  Therefore,  a  cubic 
mile  of  such  sand  would  contain  3  X  io6  cubed  or  27  X  iol8  grains, 
which  is  the  number  of  molecules  in  i  c.c.  of  gas  at  standard 
pressure  and  temperature.  Since  the  number  of  molecules  per 
cubic  centimeter  has  been  determined  by  several  independent 
methods  which  give  closely  agreeing  results,  we  may  safely 
accept  the  value  given  above  as  being  correct  within  i  per  cent. 

196.  Some  Further  Conclusions. — It  may  now  be  of  some 
interest  to  note  a  few  additional  conclusions  that  have  been 
reached  in  the  study  of  gases.    Let  us  illustrate  by  means  of  the 
gas  oxygen.      We  know  the  weight  of  i  c.c.  of  oxygen  and  the 
number  of  molecules  in  i  c.c.  at  standard  conditions;  dividing 
the  first  by  the  second  gives  the  weight  of  a  single  molecule  of 
this  gas;  this  comes  out  5 . 3  X  io~23  gram.    The  size  of  a  mole- 
cule has  also  been  approximately  determined  and  in  the  case  of 
oxygen  it  turns  out  that  the  diameter  of  a  molecule  is  approxi- 
mately 2 . 5  X  io~8  cm.    We  have  already  seen  that  the  average 

distance  between  molecules  is  about of  a  cm.  =  i .  ^  X 

3,000,000 

io~7  cm.  We  may  now  ask:  How  far,  on  the  average,  will  a 
molecule  travel  in  a  straight  line  before  it  strikes  another  mole- 
cule? This  result  can  be  calculated  when  the  diameter  and 
average  distance  apart  of  the  molecules  are  known,  and  is  called 
the  free  path ;  for  oxygen  it  is  i .  3  X  io~s  cm.  Thus  we  see  that, 
on  the  average,  a  molecule,  after  one  collision  will  travel  about 
40  times  (i.3Xio~5-i-3.3Xio~7  =  4o)  the  average  distance 
between  two  neighboring  molecules  before  striking  a  second 
molecule.  This  is  not  surprising  when  we  note  that  the  average 
distance  between  molecules  is  about  13  times  their  diameters 
(3. 3X10-74-2. 5Xio-8  =  i3). 

197.  The  Velocity  of  Molecular  Motion. — The  average  ve- 
locity with  which  molecules  travel  between  collisions  can  be  calcu- 
lated with  a  high  degree  of  certainty.    The  velocity  varies  with 
the  mass  of  the  molecule  and  its  temperature  but  is  independent 
of  the  pressure.    Molecules  of  equal  masses  have  equal  velocities 


n6  Introduction  to  General  Chemistry 

at  the  same  temperatures,  while  for  those  with  different  masses 
the  velocities  are  inversely  proportional  to  the  square  root  of  the 
mass.  At  o°  the  velocity  of  the  oxygen  molecules  is  io4  cm.  per 
second,  or  about  15  miles  per  minute.  But  since  the  free  path 
of  a  molecule  of  oxygen  is  only  i .  3  X  io~s  cm.,  it  will  experience 
many  thousands  of  collisions  in  progressing  i  cm.  At  each 
collision  its  direction  of  travel  will  change  so  that  its  actual 
progress  from  a  given  position  is  far  slower  than  its  high  velocity 
would  indicate  if  no  collisions  occurred. 

198.  The  Liquid  State. — As  a  gas  is  compressed  at  constant 
temperature  its  molecules  are  brought  closer  together,  but  other- 
wise conditions  remain  nearly  unchanged.    The  mass,  diameter, 
and  velocity  of  each  molecule  will  not  be  altered;    only  the 
average  distances  between  the  molecules  and  their  free  paths 
will  be  shortened.     It  seems  probable,  in  fact,  that  the  average 
kinetic  energy  of  a  molecule,  which  is  equal  to  one-half  the 
product  of  its  mass  and  the  square  of  its  velocity  (%mv2),  remains 
unchanged,  however  much  the  gas  is  compressed.     If  we  accept 
this  view,  we  may  easily  extend  it  to  cover  the  liquid  state,  in 
which  we  may  imagine  that  the  molecules  have  the  same  veloci- 
ties and  therefore  the  same  kinetic  energies  as  the  molecules 
of  the  vapor  of  the  liquid  have  at  the  same  temperature,  but  that 
the  crowding  of  the  molecules  is  so  great  that  their  free  paths 
are  short  compared  with  their  diameters.    However,  we  may 
think  of  the  molecules  as  able  to  progress  slowly  from  one  place 
to  another,  although  the  motion  will  be  very  irregular,  like  that  of 
persons  moving  about  in  a  dense  crowd. 

199.  Vaporization  of  a  Liquid. — It  has  already  been  stated 
that  all  of  the  molecules  of  a  given  gas  cannot  have  equal  veloci- 
ties; nor  can  a  given  molecule  always  have  the  same  velocity, 
since  at  every  one  of  the  frequent  collisions  the  velocity  will  be 
changed.    It  is  only  the  average  velocity  of  all  the  molecules 
that  remains  unchanged  as  long  as  the  temperature  remains 
constant.    The  velocities  of  the  molecules  of  a  liquid  also  are 
not  all  the  same  at  a  given  instant;  some  will  be  moving  much 
slower,  others  much  faster,  than  the  average.    If  a  fast-moving 
molecule  approaches  the  free  surface  of  the  liquid,  it  may  escape 


Molecular  Hypothesis  117 

into  the  space  above  the  liquid,  whereas  a  slow-moving  molecule, 
under  the  same  conditions,  might  not  be  able  to  escape.  Now 
the  passage  of  molecules  from  the  liquid  to  the  space  above  it  is 
nothing  but  the  evaporation  of  the  liquid.  Moreover,  we  see 
that  the  rate  of  escape  of  the  molecules,  and  therefore  the  rate 
of  evaporation,  will  be  greater  in  proportion  as  the  average 
velocity  of  the  molecules  is  increased.  Since  molecular  velocity 
increases  with  rise  of  temperature,  we  get  in  this  way  a  simple 
explanation  as  to  why  heating  a  liquid  hastens  its  evaporation. 
When  the  evaporation  of  a  liquid  goes  on  with  a  poor  supply  of 
heat,  as  when  water  evaporates  in  an  open  vessel,  the  liquid 
becomes  cooler.  Obviously  this  is  due  to  the  lowering  of  the 
average  velocity  of  the  molecules  of  the  liquid  because  of  the 
escape  of  the  faster-moving  ones. 

200.  Vapor  Pressure. — If  a  liquid  is  placed  in  a  closed  vessel 
which  it  does  not  completely  fill,  it  will  evaporate  only  until  the 
pressure  exerted  by  the  vapor  attains  a  certain  value  which  is 
definitely  determined  by  the  temperature.  For  example,  at 
20°  the  vapor  pressure  of  water  is  equal  to  that  exerted  by 
17.4  mm.  of  mercury ;  at  2  5°  it  equals  23.6  mm.  Does  the  water 
cease  to  pass  into  vapor  when  these  pressures  are  reached  ?  If 
so,  does  this  mean  that  molecules  of  water  no  longer  pass  from 
the  liquid  to  the  space  above  it?  This  would  seem  strange. 
Let  us  look  at  the  question  from  another  point  of  view.  Suppose 
we  have  a  vessel  full. of  steam  and  allow  it  to  cool.  We  know 
that  most  of  the  steam  will  condense;  only  a  little  will  remain  as 
vapor.  If  we  try  to  picture  how  this  occurs,  we  must  think  of 
some  of  the  molecules  of  vapor,  that  is,  gaseous  water,  coming 
together  first  to  form  liquid  droplets;  these  fall  to  the  bottom 
and  soon  form  a  layer  of  liquid;  other  molecules  then  strike 
this  liquid  and  remain  as  a  part  of  it.  Finally,  when  the  tempera- 
ture of  the  room,  say  20°,  has  been  reached,  the  pressure  within 
the  vessel  will  have  fallen  to  17.4  mm.  of  mercury,  and  most  of 
the  water,  but  not  all,  will  have  condensed  to  the  liquid  state. 
It  is  important  to  note  that  at  a  given  temperature,  say  20°,  the 
same  final  vapor  pressure  is  reached  whether  steam  condenses 
or  water  evaporates. 


1 1 8  Introduction  to  General  Chemistry 

201.  Equilibrium  between  Liquid  and  Vapor. — A  very  impor- 
tant question  now  confronts  us:   Do  water  molecules  cease  to 
pass  from  the  vapor  into  the  liquid  when  at  20°  the  pressure 
reaches  17.4  mm.?     If  so,  Why?    Would  it  not  seem  more 
reasonable  to  suppose  that/or  every  molecule  that  passes  from  the 
vapor  into  the  liquid  there  is  another  that  leaves  the  liquid  and  passes 
into  the  vapor  ?     This  supposed  state  of  affairs  would  correspond 
to  that  in  which  the  number  of  customers  in  a  large  shop  remains 
substantially  constant  during  a  given  hour  of  the  day,  by  reason 
of  the  fact  that  in  each  minute  as  many  persons  enter  the  shop 
as  leave  it.     When  at  constant  temperature  the  vapor  pressure 
of  a  liquid  has  reached  a  constant  value,  we  say  there  is  equi- 
librium between  liquid  and  vapor;   and  it  would  seem  from  the 
discussion  above  that  this  condition  does  not  represent  a  state  of 
rest  or  inaction,  but  one  in  which  two  opposing  actions  exactly 
counteract  one  another. 

202.  Molecular  Attraction. — If  we  think  of  the  matter  criti- 
cally, we  may  wonder  why  molecules  of  cooling  water  vapor  col- 
lect into  drops.    Perhaps  there  is  a  sort  of  attraction  between  the 
molecules  that  holds  them  together.    If  so,  why  should  it  seem 
to  be  more  effective  at  lower  than  at  higher  temperatures?    If, 
in  reality,  one  molecule  has  some  attraction  for  another,  must 
we  suppose  that  this  attraction  increases  with  fall  of  tempera- 
ture ?    Would  it  not  be  sufficient  to  assume  a  constant  attrac- 
tion of  each  molecule  of  water  for  every  other?    Suppose,  now, 
the  vapor  of  water  is  very  hot;  then  the  molecules  will  be  moving 
with  such  great  velocities  that  if  two  of  them  collide  they  will 
rebound,  exactly  as  a  rubber  ball,   thrown  downward,  will 
rebound  on  striking  the  floor,  although  gravitational  attraction 
tends  to  keep  it  on  the  floor.    But  suppose  the  vapor  to  be 
cooled;  its  molecules  will  then  have  smaller  velocities  and  some 
may  be  moving  so  slowly  that  upon  collision  they  remain  in 
contact.     Other  slow-moving  molecules,  striking  by  accident 
a  pair  of  molecules  so  formed,  may  add  themselves  to  it,  and 
in  this  way  the  droplet  of  water  could  be  formed.    There 
are  also  other  reasons  for  assuming  that  molecules  attract  one 
another. 


Molecular  Hypothesis  119 

203.  The  Solid  State. — The  most  striking  physical  difference 
between  a  solid  and  a  liquid  is  the  rigidity  of  the  former.    This 
property  of  solids  can  most  easily  be  accounted  for  by  assuming 
that  the  molecules  are  not  free  to  move  about  as  in  the  case  of  a 
liquid,  where  the  freedom  of  motion  is  comparable  to  that  of 
people  in  a  crowd,  but  that  each  molecule  remains  in  its  place 
with  respect  to  the  whole  solid,  as  well  as  to  its  neighboring  mole- 
cules.   It  is  not  necessary  to  think  of  the  molecules  as  being 
absolutely  at  rest.     It  is  more  likely  that  each  molecule  has  a 
vibrating  motion  at  all  temperatures  above  absolute  zero  and  that, 
in  fact,  its  kinetic  energy  is  as  great  as  it  would  be  if  the  mole- 
cule were  in  the  vapor  state  at  the  same  temperature. 

204.  Crystals. — Pure  chemical  substances  in  the  solid  state 
usually  form  crystals.    The  crystals  of  a  given  substance  all  have 
the  same  general  form.    Thus,  for  example,  the  crystals  of 
common  salt,  when  perfect,  are  all  cubical  in  form,  while  those  of 
quartz  occur  as  hexagonal  prisrns.    If  we  think  of  a  crystal  as 
built  up  of  molecules,  it  is  natural  to  wonder  whether  the  mole- 
cules are  present  in  haphazard  fashion,  like  potatoes  in  a  barrel,  or 
if  they  may  not  perhaps  be  arranged  in  some  systematic  manner, 
like  bricks  in  a  wall  or  balls  in  a  regular  pile.    The  probability 
that  the  molecules  of  a  crystal  are  arranged  in  a  definite  and  regular 
manner  is  greatly  increased  when  it  is  known  that  there  are  exactly 
as  many  types  of  crystalline  form  as  there  are  possible  regular 
arrangements  of  points  in  space. 

Within  the  last  few  years  it  has  become  possible  by  means 
of  photographic  studies  made  by  the  use  of  X-rays  to  obtain  very 
precise  information  regarding  the  arrangement  of  molecules 
forming  a  crystal.  As  a  result  we  now  know  quite  definitely  the 
molecular  structure  of  a  number  of  crystals. 

205.  The  Melting  of  Crystals. — Pure  crystalline  substances 
have  definite  melting  temperatures;   thus,  ice  melts  at  o°,  and 
potassium   nitrate   at   339°.    Increase    of    temperature   must 
increase  the  intensity  of  molecular  vibration;  at  some  tempera- 
ture (the  melting-point)  this  vibration  seems  to  become  so  great 
that  the  systematic  structure  of  the  crystal  is  wrecked,  leaving 
only  an  irregularly  mixed  mass  of  molecules,  forming  the  resulting 


I2O  Introduction  to  General  Chemistry 

liquid.  Crystals  cannot  be  heated  above  their  melting-points. 
Ice,  for  example,  although  it  may  be  melting  on  the  surface,  is 
never  hotter  than  zero. 

206.  Supercooling. — On  the  other  hand,  water  may  be  cooled 
2  or  3  degrees  below  zero  without  freezing,  if  it  is  kept  quiet  and 
is  not  in  contact  with  ice.     Such  supercooled  water  immediately 
begins  to  freeze  if  touched  with  a  piece  of  ice.     This  phenomenon 
is  a  common  one  and  is  easily  explained.     In  order  that  the 
formation  of  a  crystal  can  start,  a  certain  minimum  number 
of  molecules  must  come  together  in  the  proper  positions.    But 
this  exact  arrangement  of  the  several  molecules  necessary  may 
not  readily  occur,  especially  as  immediately  above  the  freezing- 
point  (which  is  the  same  as  the  melting-temperature)  the  mole- 
cules are  vibrating  so  fast  that  they  are  just  able  to  shake  apart 
this  regular  arrangement  (that  is,  to  melt  the  crystal).    At  a 
little  lower  temperature  the  molecular  motion  is  less  and  therefore 
the  conditions  are  more  favoraffle  for  the  starting  of  crystalliza- 
tion.   However,  if  a  crystal  of  the  substance  is  present,  then 
supercooling  does  not  occur,  but  the  liquid  at  once  begins  to 
crystallize  (freeze)  at  the  temperature  of  its  melting-point.    The 
reason  is  obvious:  now  each  molecule  that  touches  the  crystal 
can  find  its  proper  lodging-place,  and  so  crystalline  growth  can 
continue. 

207.  Solutions. — In  a  solution  the  molecules  of  the  dissolved 
substance  must  be  very  uniformly  distributed;  it  would  seem, 
therefore,  that  they  may  be  moving  about  freely  among  the 
molecules  of  the  solvent,  being  carried  from  place  to  place  by 
their  own  motions.    The  process  of  dissolving  of  a  substance 
would  closely  resemble  that  of  evaporation,  and  the  crystalliza- 
tion of  a  solid  from  its  solution  would  correspond  to  the  conden- 
sation of  a  vapor  to  a  liquid.    In  fact,  we  may  imagine  that  in  the 
case  of  a  saturated  solution  in  contact  with  the  crystals  of  a 
substance  we  have  a  state  of  equilibrium  as  a  result  of  the  passage 
of  molecules  into  and  out  of  the  solution  at  exactly  equal  rates. 


CHAPTER  XI 
THE  ATOMIC  HYPOTHESIS  AND  ATOMIC  WEIGHTS 

208.  Dalton's  Atomic  Hypothesis. — The  application  of  the 
Atomic-Molecular  Hypothesis  to  the  explanation  of  chemical 
phenomena  was  first  made  by  John  Dal  ton  of  Manchester  in 
1803.    Long  before  this  time  Bernoulli  had  proposed  the  Kinetic- 
Molecular  Hypothesis  as  an  explanation  of  the  physical  behavior 
of  gases,  and  Dalton,  knowing  this  view  of  the  nature  of  matter, 
sought  to  explain  the  difference  in  solubility  in  water  of  different 
gases  as  due  to  a  possible  difference  in  size  of  their  molecules. 
But  how  could  this  imagined  difference  be  discovered?    At  this 
date,  1803,  the  theory  of  the  indestructibility  of  matter  and  the 
doctrine  of  elements  were  well  established,  owing  to  the  work  of 
Lavoisier,  a  quarter  of  a  century  earlier,  as  well  as  the  labors  of 
many  able  chemists  of  the  intervening  period.     It  was  generally 
accepted  that  the  formation  of  a  substance  was  due  to  the  union 
of  the  elements  composing  it  and  in  many  cases  the  proportions 
of  the  elements  in  a  compound  were  already  known — not  very 
accurately,  it  is  true,  but  at  least  approximately.    Dalton  wished 
to  discover  the  relative  weights  of  the  ultimate  particles  of  gases; 
but  in  order  to  do  this  he  would  have  to  know,  in  the  case  of 
hydrogen  and  oxygen,  for  example,  in  addition  to  knowing  the 
weight  of  oxygen  that  would  combine  with  a  given  weight  of 
hydrogen,  the  relative  numbers  of  ultimate  particles  of  the  two 
gases  that  combine  with  one  another  in  the  formation  of  water. 
As  Dalton  had  no  experimental  means  of  discovering  the  informa- 
tion he  lacked,  he  simply  assumed  that  one  ultimate  particle 
of  hydrogen  united  with  one  ultimate  particle  of  oxygen  to  give 
one  ultimate  particle  of  water,  meaning  by  the  expression  ulti- 
mate particle  essentially  the  same  as  the  Greeks  and  later  philos- 
ophers meant  by  the  terms  "atom"  or  "molecule,"  that  is,  the 
smallest  possible  particle  of  the  substance. 

209.  Finding  the  Relative  Weights  of  Atoms. — Now,  since  ig. 
of  hydrogen  unites  with  8  g.  of  oxygen  to  form  9  g.  of  water, 


122  Introduction  to  General  Chemistry 

the  ultimate  particle  or  atom  of  oxygen  must  weigh  eight  times 
as  much  as  the  ultimate  particle  or  atom  of  hydrogen;  and  the 
ultimate  particle  of  water,  in  this  case  the  molecule,  must  weigh 
nine  times  as  much  as  an  atom  of  hydrogen.  In  making  such 
suppositions  Dal  ton  also  assumed  that  all  of  the  atoms  of  hydro- 
gen were  exactly  alike  in  size,  weight,  and  all  other  properties ; 
that  each  atom  of  oxygen  was  exactly  like  every  other  atom  of 
this  element,  but  entirely  different  from  an  atom  of  any  other 
element.  Dalton  knew  that  the  same  pair  of  elements  often 
form  two  or  more  compounds  in  which  the  constituents  are 
present  in  different  proportions.  This  forced  him  to  assume 
also  that  in  such  cases  the  atoms  unite,  not  only  one  to  one,  but 
also  one  to  two,  or  one  to  three,  etc.  In  order  that  the  student 
may  have  a  perfectly  clear  notion  of  the  matter,  we  may  sum- 
marize by  stating  that  Dalton  assumed  that  a  molecule  of  water 
is  composed  of  one  atom  of  hydrogen  and  one  atom  of  oxygen, 
and  then  reached  the  conclusion  that  an  atom  of  oxygen  was 
eight  times,  and  a  molecule  of  water  nine  times,  as  heavy  as  an 
atom  of  hydrogen.  But  Dalton  did  not  know,  as  we  can  see 
clearly,  whether  one  atom  of  hydrogen  unites  with  one  atom  of 
oxygen  or  with  two  or  three  of  oxygen,  or  whether  two  or  perhaps 
three  atoms  of  hydrogen  unite  with  one  of  oxygen  to  form  a 
molecule  of  water:  it  was  all  a  guess.  But  it  must  also  be  clear 
that  if  we  could  discover  the  numbers  of  atoms  of  hydrogen  and 
oxygen  in  a  molecule  of  water  we  could  find  the  relative  weights  of 
the  two  atoms,  knowing  the  percentages  of  hydrogen  and  oxygen 
in  water.  Now  the  question  is :  How  can  we  discover  the  num- 
ber of  atoms  of  each  kind  in  a  molecule  of  a  substance  ? 

210.  The  Application  of  Avogadro's  Hypothesis. — Suppose 
we  accept  Avogadro's  suggestion  that  equal  volumes  of  all  gases 
at  the  same  temperature  and  pressure  contain  the  same  number 
of  molecules,  and  see  to  what  conclusion  we  are  led.  Let  us 
represent  by  N  the  number  of  molecules  in  22.4  liters  of  any 
gas  under  standard  conditions.  Now  according  to  Dalton's 
suggestion  one  molecule  of  a  given  substance  will  contain  one, 
two,  three,  or  some  small  whole  number  of  atoms  of  a  given 
element,  but  cannot,  by  reason  of  the  assumed  indivisible  nature 


Atomic  Hypothesis  and  Atomic  Weights  123 

of  an  atom,  contain  a  fraction  of  an  atom.  Let  us  consider  the 
gas  ammonia  as  an  example.  Ammonia  is  composed  of  17.8 
per  cent  of  hydrogen  and  82.2  per  cent  of  nitrogen,  and  nothing 
else.  One  molecule  of  ammonia,  according  to  Dalton's  sug- 
gestion, contains  one,  two,  three,  or  four,  or  at  least  some  small 
number  of  atoms  of  hydrogen.  Now,  if  22.4  liters  of  ammonia 
gas  under  standard  conditions  contain  N  molecules,  then  this 
volume  of  the  gas  must  contain  iXTV,  2XN,  3X7V,  or  some 
small  number  of  times  N  atoms.  The  least  number  of  hydrogen 
atoms  that  could  possibly  be  contained  in  2  2 . 4  liters  of  ammonia 
is  N,  but  the  true  number  may  be  2  XN,  which  we  may  write  2  TV, 
or  it  may  be  greater,  as  37V  or  47V,  so  far  as  we  know;  only  it 
must  be  N  or  some  small  whole  number  of  times  N  if  we  assume 
that  there  are  N  molecules  in  22.4  liters  of  the  gas  and  also 
assume,  with  Dal  ton,  that  each  molecule  of  the  gas  contains 
one,  two,  three,  or  some  small  number  of  atoms  of  hydrogen.  In 
the  case  of  any  other  gaseous  compound  of  hydrogen  we  should 
conclude,  according  to  Avogadro,  that  22.4  liters  of  the  gas 
contained  N  molecules  and  that  each  molecule  contained  one, 
two,  three,  or  some  other  small  number  of  hydrogen  atoms,  the 
smallest  possible  number  being  one  atom  of  hydrogen  to  the 
molecule,  and  therefore  that  22.4  liters  of  the  gas  would  contain 
TV,  2N,  or  $N,  etc.,  atoms  of  combined  hydrogen. 

211.  The  Number  of  Atoms  and  Weight  of  Hydrogen  in  22.4 
Liters. — According  to  Dal  ton  all  hydrogen  atoms  are  alike  and 
each  has  a  definite  weight,  so  that  the  weight  of  N  atoms  of 
hydrogen  would  be  a  perfectly  definite  weight  of  this  element. 
The  weight  of  2N  atoms  of  hydrogen  would,  of  course,  be  twice 
that  of  N  atoms,  etc.  It  seems  reasonable  to  think  that  it  would 
be  likely  to  happen  that  in  some  of  the  gaseous  compounds  of 
hydrogen  the  molecules  would  contain  but  one  atom  of  hydrogen 
each.  In  such  a  case  22.4  liters  would  contain  N  atoms  of 
combined  hydrogen,  having  a  definite  weight.  Now  as  such 
gases  contain  the  minimum  possible  number  of  atoms  of  hydro- 
gen in  each  molecule,  namely,  one,  and  as  we  assume  that  all 
gases  contain  N  molecules  in  22.4  liters,  then  such  gases  would 
contain  the  minimum  possible  weight  of  hydrogen  in  this  volume. 


124  Introduction  to  General  Chemistry 

As  a  matter  of  fact  we  actually  find  that  in  22 . 4  liters  of  the  vari- 
ous gaseous  compounds  of  hydrogen  the  weight  of  this  element  is 
in  no  case  less  than  i  g.  In  this  volume  of  any  definite  gas  there  is 
either  no  combined  hydrogen  or  there  is  at  least  i  g.:  the  minimum 
weight  of  hydrogen  is  i  g. 

212.  The  Explanation  of  the  Laws  of  Minimum  and  Multiple 
Weights. — In  other  gaseous  compounds  of  hydrogen  we  find  in 
22.4  liters  larger  weights  of  combined  hydrogen,  but  these 
weights  are  then  either  2  g.,  3  g.,  or  some  whole  multiple  of  the 
minimum  weight.     Is  it  not  logical  then  to  think  that  in  such 
gases  as  hydrogen  chloride,  where  the  minimum  weight,  i  g., 
of  combined  hydrogen  is  contained  in  2  2 . 4  liters  of  the  compound 
gas,  the  molecule  contains  but  one  atom  of  hydrogen,  and  that 
in  acetylene,  where  2  g.  of  hydrogen  are  found  in  22.4  liters 
each  molecule  contains  two  atoms  of  hydrogen,  while  in  methane 
with  4  g.  of  hydrogen  in  the  same  volume,  there  are  four  atoms 
of  hydrogen  per  molecule?    Undoubtedly  so.    We  see  then 
that  we  have  in  the  assumptions  made  by  Avogadro  and  Dal  ton 
the  basis  of  an  explanation  of  the  remarkable  Laws  of  Minimum 
and  Multiple  Weights,  which  have  been  discovered  by  experi- 
ment; and,  because  of  the  agreement  between  theory  and  fact, 
we  are  inclined  to  think  that  perhaps  the  views  of  Avogadro  and 
Dal  ton  are  correct.     In  any  case  we  cannot  fail  to  see  that  these 
hypotheses  are  useful,  and  that,  indeed,  is  the  criterion  by  which 
the  worth  of  any  hypothesis  should  be  judged. 

213.  Application  of  the  Explanation  to  Other  Elements. — The 
question  now  arises  whether  the  simple  explanations  of  the 
laws  of  minimum  and  multiple  weights  may  be  applied  to  ele- 
ments other  than  hydrogen,  and  a  very  little  thought  will  show 
that  this  must  be  the  case.    The  minimum  number  of  atoms 
of  any  given  element  in  22.4  liters  of  any  of  its  gaseous  com- 
pounds is  again  A7",  the  number  found  in  the  case  of  those  gases 
the  molecules  of  which  contain  but  one  atom  each  of  the  given 
element.    The  minimum  weight  of  this  element  is,  of  course,  the 
weight  of  N  atoms  of  the  element.     For  example,  we  find  that 
in  22 .4  liters  of  gaseous  carbon  compounds  the  minimum  weight 
of  carbon  is  12  g.     Since  this  minimum  weight  is  found  in  the 


Atomic  Hypothesis  and  Atomic  Weights  125 

gases  carbon  dioxide  and  methane,  we  conclude  that  but  one 
atom  of  carbon  is  contained  in  a  molecule  of  each.  On  the  other 
hand,  22.4  liters  of  acetylene  contain  24  g.  of  combined  carbon, 
which  is  twice  the  minimum  weight,  from  which  we  conclude 
that  in  a  molecule  of  this  gas  there  are  two  atoms  of  carbon. 
As  we  have  now  reached  the  conclusion  that  a  molecule  of  methane 
contains  four  atoms  of  hydrogen  and  one  of  carbon  we  see  that  we 
have  developed  a  method  whereby  we  can  solve  the  problem 
first  suggested  by  Dalton,  that  of  discovering  the  number  of  atoms 
of  each  sort  in  a  molecule  of  a  given  substance  at  least  in  the  case 
where  the  substance  is  a  gas,  since  it  is  evident  that  the  method 
used  for  methane  is  applicable  to  any  gaseous  substance. 

214.  The  Number  of  Atoms  of  Each  Kind  in  a  Molecule. — To 
illustrate  by  further  examples,  we  may  consider  the  cases  of  a 
few  of  the  gases  we  have  already  studied.    We  see,  by  reference 
to  Table  IV,  that  in  22.4  liters  of  hydrogen  chloride  there  are 
found  the  minimum  weights  of  both  hydrogen  and  chlorine  and 
conclude  that  the  molecule  of  this  gas  is  made  up  of  one  atom 
each  of  hydrogen  and  chlorine.     We  also  see  by  Table  IV  that 
the  unit  volume  of  ammonia  contains  the  minimum  weight  of 
nitrogen  and  three  times  the  minimum  weight  of  hydrogen,  and 
decide  that  in  a  molecule  of  ammonia  one  atom  of  nitrogen  must 
be  united  with  three  atoms  of  hydrogen.     In  all  other  cases  the 
reasoning  is  equally  simple,  so  that  the  student  will  have  no 
trouble  in  deciding  upon  the  number  of  atoms  of  each  kind  in  a 
molecule  of  each  of  the  gases  mentioned  in  Table  IV. 

215.  The  Number  and  Kind  of  Atoms  in  a  Molecule  Shown 
by  the  Formula. — We  are  now  in  position  to  notice  a  most  re- 
markable fact,  which  the  following  examples  will  illustrate.     One 
molecule  of  hydrogen  chloride  contains  one  atom  of  hydrogen  and 
one  of  chlorine,  and  its  formula  is  HCl;  one  molecule  of  ammonia 
contains  three  atoms  of  hydrogen  and  one  atom  of  nitrogen,  and 
its  formula  is  NH3 ;  one  molecule  of  methane  contains  four  atoms 
of  hydrogen  and  one  atom  of  carbon,  and  its  formula  is  CH4; 
in  each  case  the  number  of  atoms  of  each  element  is  the  same  as  ihe 
number  of  symbol  weights  of  that  element  in  the  formula  of  the  sub- 
stance!   And  that  the  same  thing  is  true  for  all  gases  of  Table  IV 


126  Introduction  to  General  Chemistry 

may  readily  be  found  by  considering  each  separate  case  in  the 
same  way  as  we  did  those  of  three  of  the  gases.  In  every  case, 
therefore,  the  formula  shows  not  only  the  weight  of  each  element  in 
22.4  liters  of  the  gas  but  also  the  number  of  atoms  of  each  element 
in  one  molecule  of  the  substance. 

216.  Symbol  Weights  and  Atomic  Weights. — But  we  may 
now  inquire,  Why  should  this  be  true  ?  To  answer  this  question, 
we  will  recall  that  the  minimum  weight  of  any  element  in  22 . 4 
liters  of  its  gaseous  compounds  is  the  weight  of  N  atoms  of  that 
element.  If  N  atoms  of  hydrogen  weigh  i  g.  and  N  atoms  of 
carbon  weigh  12  g.,  then  one  atom  of  carbon  must  be  12  times  as 
heavy  as  an  atom  of  hydrogen.  In  a  similar  way  we  are  led  to 
conclude  that  an  atom  of  nitrogen  is  14  times,  and  an  atom  of 
oxygen  16  times,  as  heavy  as  an  atom  of  hydrogen.  Analogous 
relations  must  likewise  exist  in  the  cases  of  all  other  elements; 
and,  therefore,  taking  the  weight  of  one  atom  of  hydrogen  as  one 
or  unity,  the  weight  of  an  atom  of  any  other  element  is  repre- 
sented by  exactly  the  same  number  as  its  symbol  weight.  For 
this  reason  a  table  of  symbol  weights  is  also  called  a  table  of 
Atomic  Weights;  and  symbol  weights  are  usually  referred  to  as 
atomic  weights.  But  we  must  remember  that  the  symbol 
weights  may  be  found  by  simple  and  direct  experiments,  inde- 
pendently of  all  suppositions  and  hypotheses,  while  atomic 
weights  are  to  be  represented  by  the  same  set  of  numbers  only 
when  we  assume  that  matter  is  made  up  of  atoms  which  unite 
in  simple  ratios  to  form  molecules  of  which  all  gases  are  assumed 
to  contain  equal  numbers  in  equal  volumes.  Briefly  stated, 
symbol  weights  are  natural  constants,  but  atomic  weights  are  the 
probable  relative  weights  of  the  atoms  of  which  we  imagine  matter 
to  be  made  up.  We  now  may  answer  the  question  proposed  in  the 
first  sentence  of  this  paragraph.  The  number  of  atoms  of  any 
sort  in  a  molecule  is  the  same  as  the  number  of  symbol  weights  of 
that  element  because  the  absolute  weight  of  an  atom  of  any  element  is 
proportional  to  its  symbol  weight.  In  this  chapter  we  have  seen 
how  the  problem  which  Dal  ton  set  for  himself  over  a  century 
ago  is  to  be  solved,  at  least  as  definitely  as  chemists  know,  up 
to  the  present  time,  how  to  solve  it.  The  key  to  the  solution  was 


Atomic  Hypothesis  and  Atomic  Weights 

the  hypothesis  of  Avogadro,  which  was  suggested  in  1811,  only 
three  years  after  Dalton's  views  first  appeared  in  print,  and 
which  was  rejected  by  Dal  ton  himself,  and  was  only  accepted 
by  the  chemical  world  at  large  half  a  century  later. 

217.  Formula  Weights  and  Molecular  Weights. — It  is  of 
course  obvious  that  the  weight  of  a  molecule  may  also  be 
expressed  in  terms  of  the  weight  of  one  atom  of  hydrogen  which 
is  taken  as  unity.    For  example,  if  one  atom  each  of  hydrogen 
and  chlorine  compose  a  molecule  of  hydrogen  chloride  and  if,  as 
we  have  seen,  an  atom  of  chlorine  weighs  35.5  times  as  much  as 
an  atom  of  hydrogen,  then  a  molecule  of  the  compound  must 
weigh  36. 5  times  as  much  as  an  atom  of  hydrogen;  and  we  say 
therefore  that  the  Molecular  Weight  of  hydrogen  chloride  is 
36.5.     The  molecular  weight  of  a  gas  has  consequently  the  same 
numerical  value  as  its  formula  weight,  the  weight  of  an  atom  of 
hydrogen  being  in  all  cases  taken  as  unity.    The  conclusion 
that  the  relative  weights  of  the  molecules  of  gases  are  propor- 
tional to  their  respective  formula  weights  follows  at  once  from  the 
assumption  of  Avogadro's  hypothesis.    But  we  now  see  also  that 
the  weights  of  gaseous  molecules,  briefly  their  molecular  weights, 
are  all  represented  by  the  same  numbers  as  their  formula  weights 
if  we  choose  the  atomic  weight  of  hydrogen  as  unity;   for  this 
reason  it  seemed  logical  to  discuss  atomic  weights  before  molec- 
ular weights.    It  is  also  evident  that  the  molecular  weight  of  a 
substance  must  be  equal  to  the  sum  of  the  atomic  weights  indicated 
by  the  formula. 

218.  The  Formulae  of  Some  Elementary  Gases. — We  are  now 
in  position  to  consider  the  meaning  of  the  fact  that  the  formulae 
H2,  O2,  N2,  and  C12  were  found  for  the  four  elementary  gases 
studied.     If  we  accept  Avogadro's  hypothesis  for  these  as  well 
as  for  compound  gases,  then  the  unit  volume  of  each  gas  must 
also  contain  N  molecules.     But  we  know  also  that  the  weight 
in  each  case  is  twice  the  minimum  weight  of  N  atoms;  for  this 
reason  we  are  forced  to  conclude  that  the  unit  volume  of  each 
gas  contains  2N  atoms,  and  hence  that  each  molecule  contains 
two  atoms.    The  same  rule  that  applies  to  compound  gases 
applies  here  also:  the  number  of  atoms  of  each  element  in  any 


128  Introduction  to  General  Chemistry 

molecule  is  the  same  as  the  number  of  symbol  weights  of  that 
element  in  the  formula  of  the  substance.  The  obvious  meaning 
of  these  facts  is  that  an  atom  of  hydrogen,  for  example,  can 
unite  with  another  atom  of  hydrogen  as  well  as  with  one  of 
chlorine  or  some  other  element.  There  remains  the  question: 
Do  we  have  any  other  evidence  of  the  truth  of  this  conclusion  ? 
Let  us  see. 

219.  The  Union  of  Hydrogen  and  Chlorine  by  Volume. — It 
will  be  recalled  that  when  hydrogen  and  chlorine  gases  unite  to 
form  gaseous  hydrogen  chloride,  one  volume  of  each  of  the 
elementary  gases  combines  to  give  two  volumes  of  the  product. 
Now  in  two  unit  volumes  of  2  2 . 4  liters  each  of  hydrogen  chloride 
there  must  be  2  N  atoms  of  hydrogen  and  2  A7  atoms  of  chlorine, 
since  we  cannot  have  less  than  one  atom  of  each  element  in  a 
molecule  of  the  compound;  but  the  two  unit  volumes  of  hydrogen 
chloride  are  formed  from  one  unit  volume  of  hydrogen  and  one  of 
chlorine  each  containing  N  molecules;  and  again  we  are  led  to 
the  conclusion  that  N  molecules  of  hydrogen  or  chlorine  contain 
2  AT  atoms  in  each  case,  or  that  each  molecule  of  either  elemen- 
tary gas  contains  two  atoms. 

220.  Gay  Lussac's  Law  of  Combining  Volumes. — The  simple 
volumetric  relation  between  hydrogen,  chlorine,  and  hydrogen 
chloride,  1:1:2,  is  not  an  exceptional  case;    other  gases  also 
exhibit  similar  simple  relations.     Thus,  two  volumes  of  hydrogen 
and  one  volume  of  oxygen  unite  and,  if  the  temperature  at  which 
the  experiment  is  carried  out  is  so  high  that  the  water  remains  in 
the  form  of  steam,  the  latter  measures  two  volumes;  so  that  the 
volume  relations  are  2:1:2.     In  the  burning  of  methane  one 
volume  of  the  gas  requires  two  volumes  of  oxygen  and  gives 
one  volume  of  carbon  dioxide  and  two  volumes  of  steam,  the 
measurements  all  being  made  at  a  sufficiently  high  temperature 
in  this  case  also  to  keep  the  steam  from  condensing.     Or,  when 
ammonia  gas  is  decomposed,  as  it  may  be  by  means  of  electric 
sparks,  two  volumes  of  ammonia  yield  one  volume  of  nitrogen 
and  three  volumes  of  hydrogen.    The  fact  that  gases  and  vapors 
of  volatile  substances  always  react  in  simple  ratios  by  volume 
was  discovered  by  Gay  Lussac  in  1808,  and  is  known  as  Gay 
Lussac's  Law  of  Combining  Volumes. 


Atomic  Hypothesis  and  Atomic  Weights  129 

221.  Explanation  of  the  Law  of  Combining  Volumes. — The 

explanation  of  this  law  will  appear  if  we  write  the  equations  of 
the  reactions  mentioned : 

H2+C12->2HC1  CH4+202->C02+2H2O 

I  Vol.   I  VOl.   2  Vols.  I  VOl.   2  Vols.   I  Vol.   2  Vols. 

2H2+O2^2H2O  2NH3->3H2+N2 

2  VOls.    I  Vol.    2  Vols.  2  Vols.  3  Vols.    I  Vol. 

Again  we  see,  as  we  did  earlier  (76,  77),  that  the  volumes  are 
the  same  as  the  coefficients  of  the  formulae  in  the  equations,  and 
this  for  the  fundamental  reason  that  one  formula  weight  always 
represents  one  unit  volume,  in  the  case  of  a  gaseous  or  volatile 
substance.  Moreover,  we  now  understand,  to  cite  the  first 
example,  that  one  unit  volume  of  hydrogen  containing  N  mole- 
cules and  2  N  atoms  will  require  2  N  atoms  of  chlorine  or  N  mole- 
cules, which,  according  to  Avogadro's  hypothesis,  will  be  found 
in  one  unit  volume  of  chlorine.  The*  reaction  will  then  produce 
2  N  molecules  of  hydrogen  chloride,  which  according  to  the  same 
hypothesis  will  occupy  two  unit  volumes.  Similar  reasoning 
may  be  applied  to  all  other  cases. 

222.  The  Degree  of  Accuracy  of  Symbol  Weights. — Before 
leaving  the  discussion  of  symbol  and  atomic  weights  we  must 
consider  the  degree  of  accuracy  of  the  statements  of  numerical 
results  made  in  early  chapters  and  summarized  in  Table  IV.     It 
is  perhaps  needless  to  point  out  that  statements  of  lengths,  areas, 
volumes,  weights,  etc.,  whether  they  refer  to  scientific  or  other 
matters,  are  in  general  more  or  less  approximate,  the  degree  of 
accuracy  aimed  at  being  determined  by  the  requirements  of  the 
case.    Thus,  if  a  stranger  in  the  city  inquires  the  distance  from 
the  City  Hall  to  the  University  and  is  told  by  a  policeman  that 
it  is  seven  miles,  the  answer  is  quite  as  accurate  as  necessary. 
But  such  approximate  statements  of  distance  would  not  satisfy 
the  requirements  of  a  surveyor  who  wished  to  make  an  accurate 
map  of  the  city.    Up  to  about  twenty-five  years  ago  the  most 
accurate  analyses  of  water  indicated  that  2  g.  of  hydrogen  were 
combined  with  15.96  g.  of  oxygen.    As  all  chemists  know  that 
in  every  analysis  there  is  inevitably  some  experimental  error 


130  Introduction  to  General  Chemistry 

of  greater  or  less  magnitude,  it  was  thought  that  the  true  weight 
of  oxygen  combined  with  exactly  2  g.  of  hydrogen  was  exactly 
1 6  g.  It  then  became  apparent  from  the  new  researches  of  a 
number  of  chemists  that  the  error  in  the  accepted  results  was 
greater  than  suspected,  and,  moreover,  that  the  true  proportion 
of  oxygen  in  water  was  less  instead  of  greater  than  the  value 
found  earlier,  the  new  experiments  leading  to  a  ratio  of  2  to  1 5 . 88, 
with  a  probable  error  of  less  than  o.oi  g.  in  the  weight  of  oxygen 
combined  with  exactly  2  g.  of  hydrogen. 

223.  O  =  i6.ooo,  the  Real  Basis  for  Symbol  and  Atomic 
Weights. — An  annoying  difficulty  now  arose  from  the  fact  that 
far  more  symbol  weights  had  been  found  by  the  analysis  of  oxygen 
compounds  than  by  the  analysis  of  compounds  with  hydrogen, 
owing  to  the  greater  accuracy  with  which  the  former  analyses 
could  be  made;  so  that  it  then  became  necessary  for  chemists  to 
decide  whether  they  should  change  the  symbol  weights  of  oxygen 
and  all  elements  whose  symbol  weights  had  been  found  by  the 
analysis  of  their  oxygen  compounds,  or  whether  they  should 
change  the  symbol  weights  of  hydrogen  and  a  few  other  elements. 
After  much  debate  the  former  policy  was  adopted  and  the  symbol 
weight  of  oxygen,  O  =  16.000,  kept  unchanged,  although  this 
made  it  necessary  to  change  the  symbol  weight  of  hydrogen  to 
i. 008.     Our  most  accurate  knowledge  of  the  composition  of 
water  is  expressed  by  the  statement  that  2X1. 008  g.  of  hydrogen 
are  combined  with  16.000  g.  of  oxygen  in  18.01 6  g.  of  water,  a 
fact  which  is  also  expressed  by  the  formula  H20,  when  we  consider 
that  H  =  i.oo8  g.  of  hydrogen  and  O  =  16.000  g.  of  oxygen. 
Oxygen  with  a  symbol  weight  of  16.000  has  thus  become  the  real 
basis  of  the  system  of  symbol  and  atomic  weights  rather  than  hydrogen 
with  a  symbol  weight  of  unity. 

224.  The  Method  of  Finding  Symbol  Weights. — The  symbol 
weights,  and  therefore  also  the  atomic  weights,  of  all  other  ele- 
ments are  now  based  upon  that  of  oxygen  taken  as  16.000;  but 
we  see  by  a  comparison  of  .the  values  given  in  a  table  of  exact 
atomic  weights  that  in  no  case  does  the  exact  value  based  on 
0- 16.000  differ  greatly  from  the  approximate  value  we  have 
previously  used.     Just  as  more  accurate  analyses  led  to  a  change 


Atomic  Hypothesis  and  Atomic  Weights  131 

in  the  symbol  weight  of  hydrogen,  so  also  newer  analyses  have 
led  and  will  continue  in  the  future  to  lead  to  a  more  exact 
knowledge  of  the  symbol  weights  of  other  elements.  We  do  not 
expect,  however,  that  the  values  accepted  at  present  for  the 
commoner  elements  will  be  changed  by  more  than  a  few  units 
in  the  second  decimal  place.  Concisely  stated,  the  matter  stands 
thus:  Approximate  symbol  weights  are  found  in  the  manner 
described  in  chap,  v,  while  the  more  exact  values  are  fixed  by  the 
most  painstaking  analyses  and  syntheses,  being  computed  on  the 
basis  of  O  =  i6 .  ooo. 

225.  Inexactness  of  the  Gas  Laws. — The  gas  laws  of  Boyle, 
Charles,  and  Avogadro  are  also  only  closely  approximate  state- 
ments of  the  facts.     For  example,  if  the  pressure  on  1,000  c.c. 
of  oxygen  under  standard  conditions  be  exactly  doubled,  the 
volume  will  become  499.3  c.c.  instead  of  exactly  500,  as  Boyle's 
law  would  indicate.     The  deviations  from  the  simple  laws  are 
thought  to  be  due  to  attractions  between  the  molecules,  on  the 
one  hand,  tending  to  diminish  the  volume,  and,  on  the  other 
hand,  to  the  fact  that  part  of  the  space  occupied  by  the  gas  is 
filled  with  the  molecules  themselves,  so  that  the  free  space  is 
reduced  to  less  than  half  if  the  volume  of  the  gas  is,  by  increase 
of  pressure,  reduced  to  half.     The  actual  deviations  from  the 
simple  law,  PF  =  a  constant,  become  negligible  if  gases  are  under 
low  pressures.     Then  the  three  great  laws  express  almost  exactly 
the  behavior  of  all  gases.     In  other  words,  if  the  barometric 
pressure  at  sea-level  were  o.oi  of  its  actual  value,  so  that  our 
standard  of  atmospheric  pressure  would  be  o.  76  cm.  of  mercury 
instead  of  76  cm.,  then  we  should  find  that  not  only  would  the 
laws  of  Boyle  and  Charles  express  with  a  high  degree  of  accuracy 
the  behavior  of  gases  under  pressures  of  this  order  of  magnitude, 
but  that  for  all  gases  the  law  of  Avogadro  would  also  hold  good 
with  as  great  a  degree  of  accuracy  as  experiment  would  enable 
us  to  determine. 

226.  Exactness  of  Avogadro's  Law  for  Corrected  Gas  Vol- 
umes.— Now,  instead  of  trying  to  weigh  and  measure  gases  under 
such  low  pressures  in  attempts  to  study  them  more  accurately, 
chemists  have  worked  at  ordinary  pressures  and  then  corrected 


132  Introduction  to  General  Chemistry 

the  data  so  obtained  so  as  to  give  the  results  that  would  theoreti- 
cally have  been  obtained  for  the  weights  of  i  liter  if  the  measure- 
ments had  been  made  at  very  low  pressures  and  the  calculations 
made  for  a  pressure  of  76  cm.  exactly  according  to  Boyle's  law. 
Working  in  this  way,  it  was  found  that  the  corrected  volume  of 
32  g.  of  oxygen,  the  weight  represented  by  O2,  is  22.41  liters 
at  o°.  It  was  then  discovered  that  exactly  this  (corrected) 
volume  of  any  other  gas  at  o°  contains,  as  nearly  as  the  deter- 
minations could  be  made,  just  the  weight  of  the  gas  which  its 
formula  indicates,  this  weight  being  calculated  from  the  most 
exact  symbol  weights.  In  others,  Avogadro's  law  would  hold 
exactly  at  low  pressures  or  also  at  ordinary  pressure  if  the  attrac- 
tions of  the  molecules  for  each  other  did  not  exist,  and  if  their 
own  volumes  were  negligible  as  compared  with  the  total  space 
occupied  by  the  gas. 

227.  A  Little  Explanation  and  Advice. — It  is  not  necessary  nor 
desirable  that  the  beginner  in  chemistry  should  pay  much  atten- 
tion to  the  matters  discussed  in  the  three  preceding  paragraphs. 
The  approximate  symbol  weights  and  the  gas  laws  in  their  sim- 
plest forms  are  sufficiently  exact  for  his  use.     It  is  much  better 
that  he  should  see  clearly  the  general  fundamental  principles 
than  that  he  should  be  perplexed  and  confused  by  the  details 
and  refinements  that  are  of  importance  only  to  the  specialist. 
If  the  beginner  continues  his  study  of  chemistry  he  will  be  sure 
to  encounter  later  these  interesting  topics,  when  he  will  be  better 
able  to  appreciate  and  understand  them ;  while  if  he  should  not  go 
farther  than  the  first  course,  he  may  feel  assured  that  he  has  be- 
come acquainted  with  the  principles  of  most  fundamental  impor- 
tance.   These  matters  are  discussed  here  in  order  to  explain  why 
the  symbol  or  atomic  weights  given  in  Tables  of  Atomic  Weights 
(see  inside  of  back  cover  of  this  book)  are  not  exactly  the  same 
as  those  we  have  used  in  the  earlier  chapters. 

228.  Means  of  Discovering  Symbol  Weights. — The  student 
will  doubtless  have  received  the  impression  from  the  study  of  the 
foregoing  chapters  that  we  can  discover  the  approximate  symbol 
weight  of  any  element  by  finding  the  minimum  weight  of  the 
element  in  the  unit  volume  of  its  gaseous  or  vaporized  com- 


Atomic  Hypothesis  and  Atomic  Weights 


133 


pounds;  and  this,  in  fact,  is  true  for  a  large  number  of  elements 
in  addition  to  the  five  included  in  Table  IV.  We  shall  now 
consider  some  facts  leading  to  a  knowledge  of  the  symbol  weights 
of  a  dozen  elements  other  than  the  five  already  studied.  These 
twelve  elements  all  form  volatile  compounds,  the  densities  of 
which  may  be  determined  by  making  experiments  at  sufficiently 
high  temperatures  and  then  calculating,  by  the  laws  of  Boyle 
and  Charles,  for  the  standard  conditions,  the  weight  of  the  com- 
pound in  22.4  liters.  Multiplication  of  the  weight  so  found  by 
the  percentage  of  the  element  in  question  in  the  compound  gives 
the  weight  of  the  element  in  22.4  liters  of  the  vapor,  as  recorded 
in  Table  X. 

TABLE  X 


Volatile  Compounds  of 
Various  Elements 

Weight  of 
Elements  in 
2  2.  4  Liters 

Symbol  and 
Symbol  Weight 

Specific  Heat 

Product  of 
Symbol  Weight 
and  SpecificHeat 

Antimony  trichloride.  .  .  . 
Arsenic  trichloride  
Bismuth  trichloride  
Cadmium  
Chromium  oxychloride  .  . 
Hydrogen  iodide  
Iron  carbonyl  

HQ-5 
75-4 
217.0 
114.0 

55-o 

127.7 

S3  •  2 

Sb  =120.2 
As-=   75.0 
Bi  =208.5 
Cd  =112.4 
Cr  =   52.1 
I     =127.9 
Fe  =    ss  9 

0.0503 
O  .  0830 
0  .  0303 
0-0551 
O.  II2I 
0.0541 

o  1162 

6.0 
6.2 

6-3 

6.2 

5-8 
6.9 
6  5 

Lead  chloride  

207  .  2 

Pb  =206.9 

o  .  0304 

6.3 

Mercury  

2O2  .  2 

Hg  =  200  .  o 

o  .  0308 

6.1 

Nickel  carbonyl  

59.  7 

Ni  =  58.7 

o.  1084 

6.4 

Phosphorus  trichloride  .  . 
Zinc  chloride  

31-9 
63.5 

P    =  31-0 
Zn  =  65.4 

O.  2O2O 
0.0935 

6-3 
6.1 

In  all  cases  the  compounds  are  such  as  contain  the  mini- 
mum weight  of  the  element  the  symbol  of  which  appears  in  the 
table;  that  is  to  say,  we  do  not  know  any  other  volatile  com- 
pounds in  the  respective  cases  containing  appreciably  smaller 
weights  in  the  unit  volume.  The  weights  so  found  are,  therefore, 
approximately  the  symbol  weights  in  each  case.  The  exact 
symbol  weight  in  any  case  is  then  calculated  from  the  accurately 
determined  percentage  composition  of  some  compound  of  the 
element  with  an  element  of  exactly  known  symbol  weight. 

229.  The  Product  of  Specific  Heats  and  Symbol  Weights. — 
There  are  a  great  many  elements  which  do  not  form  gaseous  com- 
pounds, or  compounds  which  are  sufficiently  volatile  without 


134  Introduction  to  General  Chemistry 

decomposition,  to  enable  us  to  find  their  symbol  weights  in  the 
manner  above  indicated.  Very  fortunately  other  methods  have 
long  been  known  by  which  the  desired  end  can  be  attained.  We 
shall  now  consider  one  of  these  methods. 

A  very  simple  relation  was  discovered  nearly  a  century 
ago,  by  Dulong  and  Petit,  between  symbol  weights  and  specific 
heats  of  solid  free  elements.  The  amount  of  heat  required  to 
raise  the  temperature  of  a  given  weight  of  iron  i°  would  raise 
the  temperature  of  an  equal  weight  of  water  only  o.  1162°;  and 
we  say,  therefore,  that  the  specific  heat  of  iron  is  o.  1162.  The 
specific  heats  of  the  other  elements  of  Table  X  are  given  in  the 
fourth  column.  If,  now,  we  multiply  the  specific  heat  of  an 
element  by  its  symbol  weight  we  get  the  remarkable  series  of 
products  contained  in  the  last  column  of  the  table,  where  we  see 
that  the  values  are  nearly  the  same  in  all  cases.  Does  it  not 
seem  probable  that  the  law  which  we  find  applying  to  the  ele- 
ments of  Table  X  would  also  hold  good  for  other  solid  elements 
even  though  they  do  not  form  easily  volatile  compounds?  If 
so,  it  is  clear  that  in  order  to  find  the  approximate  symbol  weight 
of  an  element  we  have  only  to  divide  6. 4  by  its  specific  heat,  which 
latter  constant  can  in  general  be  found  by  a  simple,  direct 
experiment.  As  a  matter  of  fact,  this  method  has  been  of  much 
service  in  just  this  way. 

230.  Interpretation  of  the  Law  of  Dulong  and  Petit. — The  law 
of  Dulong  and  Petit  is,  moreover,  of  the  greatest  interest  and 
importance  when  viewed  from  the  theoretical  standpoint.  The 
product  of  the  specific  heat  and  symbol  weight  is  obviously  the 
quantity  of  heat  required  to  raise  the  temperature  of  the  symbol 
weight  of  an  element  one  degree;  and  this  amount  of  heat  is  the 
same  for  one  element  as  for  another.  But  the  symbol  weights 
of  various  elements  are  the  weights  of  equal  numbers  of  atoms, 
and  we  see,  therefore,  that  it  requires  equal  amounts  of  heat 
to  raise  the  temperature  of  equal  numbers  of  various  kinds  of 
atoms  by  one  degree!  The  products  of  symbol  weights  and 
specific  heats  are  generally  called  Atomic  Heats;  so  that  the 
Law  of  Dulong  and  Petit  may  be  stated  thus:  The  atomic  heats 
of  the  solid  elements  are  equal. 


CHAPTER  XII 

THE    HALOGENS    AND   THEIR   COMPOUNDS  WITH   HYDROGEN 

AND  METALS 

231.  The  Halogens. — The  elements  fluorine,  chlorine,  bro- 
mine, and  iodine  bear  a  close  resemblance  to  one  another  in  their 
properties  and  chemical  behavior;    collectively  they  are  called 
the  halogens  (from  halite,  the  scientific  name  for  rock  salt). 
In  the  present  chapter  we  shall  first  briefly  review  what  has 
already  been  learned  about  chlorine  and  some  of  its  compounds, 
and  then  after  a  more  extensive  consideration  of  the  chemistry 
of  chlorine  take  up  a  study  of  the  remaining  members  of  this 
important  group  of  elements. 

232.  Resume  of  Facts  Already  Learned. — We  know  that 
common  salt,  NaCl,  is  the  most  abundant  compound  of  chlorine ; 
it  forms  the  raw  material  from  which  all  other  compounds  of 
chlorine  as  well  as  the  free  element  are  made.    The  action  of 
sulfuric   acid   on   salt   (103)   yields  hydrochloric  acid  which, 
by  electrolysis  (43)  or  by  the  action  of  lead  dioxide,  gives  free 
chlorine  (167).    With  bases  or  metallic  oxides  hydrochloric  acid 
yields  chlorides,  as  illustrated  by  the  following  reactions: 

KOH+HC1->KC1+H2O  (107) 

MgO+  2HCl^MgCl2+H20.  (143) 

Chlorides  also  result  when  carbonates  are  treated  with  hydro- 
chloric acid  (163): 

Na2CO3+2HCl->2NaCl+CO2+H2O 
CaC03+2HCl-»CaCl2+CO2+H2O. 

It  will  be  recalled  that  the  chlorides  of  silver,  lead,  and 
univalent  mercury  are  almost  insoluble  in  water  (167, 169,  182); 
these  salts  are  easily  obtained  by  the  action  of  solutions  of  hydro- 
chloric acid  or  any  soluble  chloride  on  solutions  of  soluble  salts 
of  these  metals,  thus: 

AgNO3+HCl->AgCl+HNO3 

Pb(NO3)2+2NaCl^PbCl2-f2NaNO3. 


136  Introduction  to  General  Chemistry 

The  metals  which  react  with  hydrochloric  acid  set  free  hydrogen 
and  are  themselves  converted  into  chlorides,  for  example: 

Zn+2HCl->ZnCl2+H2  (149) 

(174) 


Chlorides  also  result  from  the  direct  union  of  chlorine  with  other 
elements  : 

H2+C12^2HC1  (44) 

2A1+3C12->2A1C13.  (174) 

233.  The  Occurrence  of  Chlorine  Compounds  in  Nature.— 

Free  chlorine  does  not  occur  in  nature.  If  free  chlorine  were 
present  in  nature  it  would  very  soon  unite  with  other  substances 
to  form  compounds.  Common  salt  is  by  far  the  most  abundant 
natural  compound  of  the  element.  It  occurs  as  a  mineral,  rock 
salt  (halite)  ,  and  as  dissolved  salt  in  sea-water  and  the  waters  of 
salt  lakes  and  springs.  Sea-water  contains  about  3  per  cent, 
while  the  water  of  Great  Salt  Lake  in  Utah  contains  about 
20  per  cent,  of  salt.  Rock  salt  has  doubtless  been  formed  in 
past  geological  times  by  the  slow,  natural  evaporation  of  sea- 
water.  Other  chlorides,  particularly  those  of  potassium,  KC1; 
magnesium,  MgCl2;  silver,  AgCl;  and  lead,  PbCl2,  are  also  found 
in  nature. 

234.  The  Discovery  of  Chlorine.  —  Free  chlorine  was  first 
made  by  the  Swedish  chemist  Scheele,  in  1774,  and  therefore 
practically  at  the  same  time  that  Lavoisier  in  France  discovered 
the  true  explanation  of  burning.    Scheele  made  chlorine  by 
the  action  of  hydrochloric  acid  on  manganese  dioxide,  a  mineral 
having  the  formula  Mn02,  and  therefore  an  oxide  of  the  metallic 
element  manganese.     The  reaction  occurs  thus: 

4HC1+  MnO2-»MnCl2-h  C12+  2H2O. 

Chlorine  was  not  thought  to  be  an  element  until  nearly  forty 
years  after  its  discovery,  but  was  believed  to  be  an  oxide  of 
hydrochloric  acid,  until  a  famous  English  chemist,  Sir  Humphrey 
Davy,  showed  by  conclusive  experiments  that  it  did  not  contain 
oxygen  and  was  really  an  elementary  substance. 


Halogens  with  Hydrogen  and  Metals 


137 


235.  The  Preparation  of  Chlorine  from  Hydrochloric  Acid.— 

We  have  already  seen  (167)  that  chlorine  is  formed  when  lead 
dioxide  is  warmed  with  hydrochloric  acid : 

4HCl+PbO2->PbCl2+Cl2+  2H2O. 

This  reaction  is  entirely  analogous  to  the  one  between  hydro- 
chloric acid  and  manganese  dioxide  mentioned  in  the  preceding 
paragraph,  and  since  the  last  substance  is  cheaper  than  lead 
dioxide  it  is  the  one  commonly  used  in  the  laboratory  for  the 
preparation  of  chlorine.  The  experimental  method  consists  in 
adding  to,  say,  100  g.  of  granular  manganese  dioxide  contained 
in  a  flask  about  300  c.c.  of  concentrated 
hydrochloric  acid  and  warming  gently : 

4HC1+  MnO2-»MnCl2+ Cl2-f-  2H2O. 

Manganese  chloride  is  an  easily  soluble 
salt  which  forms  pink  crystals  of  a 
hydrate.  MnCl2  •  4H20. 

An  excellent,  though  expensive, 
method  of  making  small  amounts  of 
chlorine  for  experimental  work  in  the 
laboratory  consists  in  allowing  concentrated  hydrochloric  acid  to 
drop  slowly  on  to  solid  potassium  permanganate,  KMn04  (Fig.  30) . 
The  latter  substance  is  one  of  the  most  powerful  oxidizing  agents 
and  reacts  rapidly  in  the  cold  with  hydrochloric  acid,  thus: 

i6HCl+  2KMnO4->2KCl+  2MnCl2+  5C12+8H2O. 

Since  the  rate  of  production  of  chlorine  is  easily  regulated  by 
control  of  the  rate  of  flow  of  the  acid,  the  method  is  a  very  con- 
venient one  for  the  lecture  table. 

236.  Chlorine,  a  Poisonous  Gas. — The  chlorine  which  is  given 
off  is  a  heavy,  yellowish,  poisonous  gas  having  an  exceedingly 
violent  action  on  all  mucous  membranes.     It  is  the  gas  which 
was  first  used  with  such  frightful  effect  in  the  trenches  in  the 
European  war.     Great  care  must  be  exercised  to  prevent  the  escape 
of  appreciable  amounts  of  chlorine  into  the  air  of  the  laboratory 


FIG.  30 


138  Introduction  to  General  Chemistry 

and  to  avoid  as  far  as  possible  inhalation  of  the  gas.  Waste 
chlorine  is  easily  absorbed  when  passed  into  a  solution  of 
caustic  soda. 

237.  The  Electrolytic  Preparation  of  Chlorine. — We  have 
already  learned  (43)  that  chlorine  is  formed  when  hydrochloric 
acid  is   electrolyzed.     By  means  of   the  Brownlee  apparatus 
shown  in  Fig.  21  it  is  found  that  equal  volumes  of  hydrogen 
and  chlorine  are  formed  when  the  concentrated  acid  is  used. 
If,  however,  very  dilute  acid  is  used,  then  the  products  are 
largely  hydrogen  and  oxygen  formed  by  the  decomposition  of 
the  water,  and  very  little  chlorine  is  set  free.    A  complete 
explanation  of  this  curious  fact  is  not  possible  until  certain 
matters  treated  in  a  following  chapter  have  been  considered; 
but  it  may  be  stated  that  hydrochloric  acid  is  more  easily 
decomposed  than  water  by  the  electric  current,  and  that  if  much 
of  the  former  is  present  in  a  water  solution  it  is  decomposed  by 
preference  to  the  water.     In  the  electrolysis  apparatus  the  poles 
or  electrodes  are  sticks  of  carbon.    The  hydrogen  is  liberated 
at  the  negative  pole,  the  chlorine  at  the  positive  pole. 

238.  The  Electrolysis  of  Common  Salt. — The  electrolysis  of 
a  concentrated  solution  of  common  salt  is  by  far  the  most  im- 
portant practical  method  for  the  manufacture  of  chlorine.     It 
is  a  process  which  is  carried  out  on  a  very  large  scale,  as  at 
Niagara  Falls,  where  electrical  power  is  cheap  and  yields  not  only 
chlorine  but  also  hydrogen  and  caustic  soda.     We  might  expect 
the  products  of  the  electrolysis  of  salt  to  be  sodium  and  chlorine, 

2NaCl-»2Na+Cl2, 

but  when  we  recall  that  sodium  reacts  at  once  with  water  to 
form  hydrogen  and  sodium  hydroxide  (caustic  soda),  the  actual 
result  appears  reasonable.  A  more  complete  explanation  must 
be  deferred  until  later.  As  in  the  case  of  the  electrolysis  of 
hydrochloric  acid  the  chlorine  is  set  free  at  the  positive  electrode, 
which  is  a  carbon  plate,  while  the  sodium  and  hydrogen  are 
formed  at  the  negative  electrode. 

239.  Deacon's  Process. — Before  the  electrical  method  just 
described  was  used  practically,  a  process  invented  by  Deacon 


Halogens  with  Hydrogen  and  Metals  139 

was  the  cheapest  technical  method  of  making  chlorine.  This 
process  is  based  on  the  fact  that  a  mixture  of  hydrogen  chloride 
gas  and  oxygen  react  at  a  high  temperature  to  form  chlorine  and 

water, 

4HC1+O2->2C12+2H2O. 

This  reaction  scarcely  takes  place  at  all  at  ordinary  temperatures, 
and  even  at  the  most  favorable  high  temperature  it  takes  place 
very  slowly.  Deacon  discovered  that  the  reaction  could,  be 
greatly  hastened  if  the  heated  mixture  of  hydrogen  chloride  and 
oxygen  were  passed  over  broken  bricks  coated  with  copper 
chloride,  CuCl2.  A  small  amount  of  this  substance  is  able  to 
promote  the  reaction  of  almost  unlimited  amounts  of  the  reacting 
gases  without  itself  being  permanently  changed  or  destroyed. 
A  substance  that  behaves  in  this  way  is  called  a  catalytic  agent. 
Catalytic  agents  of  various  sorts  are  extensively  employed  in 
chemistry.  In  the  Deacon  process  air,  which  is  essentially  a 
mixture  of  oxygen  and  nitrogen,  may  be  used  instead  of  pure 
oxygen,  which  would  be  too  expensive  for  practical  purposes. 

240.  A  Remarkable  Phenomenon:  Chemical  Equilibrium. — 
It  is  a  remarkable  fact  that  even  under  the  most  favorable  con- 
ditions the  reaction  between  hydrogen  chloride  and  oxygen  does 
not  go  to  completion,  but  stops  while  the  gaseous  mixture  still 
contains  some  of  both  of  these  gases.  The  cause  is  discovered 
when  we  find  that  steam  and  chlorine  react  at  about  400°  to  give 
some  hydrogen  chloride  and  oxygen: 

2C12+2H2O->4HC1+O2. 

This  is,  in  fact,  exactly  the  reverse  of  the  reaction  we  have  been  con- 
sidering. It  is  plain,  therefore,  that  the  failure  of  the  reaction 
between  hydrogen  chloride  and  oxygen  to  %o  to  completion  is 
due  to  the  interaction  of  the  products,  chlorine  and  water,  to 
form  again  some  of  the  first  pair  of  gases. 

If  a  mixture  of  hydrogen  chloride  and  oxygen  in  the  propor- 
tions shown  in  the  equation  is  heated  to  a  constant  temperature, 
say  400°,  a  mixture  finally  results  in  which  all  four  of  the  sub- 
stances are  present  in  definite  proportions.  A  mixture  having 
exactly  the  same  proportions  of  each  of  the  four  substances 


140 


Introduction  to  General  Chemistry 


FIG.  31 


results  if  the  starting  substances  are  chlorine  and  water,  taken 
also  in  the  proportions  indicated  by  the  equation.  In  the 
mixture  which  finally  results,  the  four  substances  are  said  to  be 
in  a  state  of  chemical  equilibrium.  The  subject  of  chemical 
equilibrium  is  a  very  important  one  which  is  to  be  studied  in 
detail  in  the  next  chapter. 

241.  The  Physical  Properties  of  Chlorine. — Chlorine  is  a 
pale-yellow  gas,  having  a  density  about  two  and  a  half  times  as 
great  as  air.     Under  standard  conditions  one  liter  weighs  3 .  22  g. 
Chlorine  is  rather  soluble  in  water,  100  c.c.  of  water  at  20°  dis- 
solving 226  c.c.  of  the  gas.     For  this  reason  the  gas  is  not  easily 

collected  over  water;  on  account 
of  its  high  density  it  is  easily 
collected  by  the  downward  dis- 
placement of  air.  If  a  water 
solution  of  chlorine  is  cooled 
nearly  to  o°,  yellow  crystalline 
chlorine  hydrate,  having  the 
formula  C12*8H20,  is  formed. 
This  hydrate  is  very  unstable  and  decomposes  slowly  at  room 
temperature  and  rapidly  at  higher  temperatures  into  chlorine 
gas  and  water. 

242.  The  Liquefaction  of  Chlorine. — A- very  interesting  and 
important  experiment  was  once  made  with  this  hydrate  by  the 
great  English  physicist  and  chemist  Faraday,  who  was  at  the 
time   assistant   to    Sir   Humphrey   Davy    (234).     Crystals   of 
chlorine  hydrate  were  sealed  up  in  one  end  of  a  bent  glass  tube, 
as  shown  in  Fig.  31;    when  the  hydrate  was  gently  warmed 
while  the  other  end  of  the  tube  was  cooled  with  ice  a  yellow 
liquid  formed  in  the  cold  end  of  the  tube.     This  liquid  proved 
to  be  liquefied  chlorine.     It  is  a  heavy,  mobile  liquid,  which  is 
easily  obtained  from  chlorine  gas  either  by  cooling  the  latter 
to  about  40°  below  zero  at  atmospheric  pressure,  or  by  com- 
pressing it  to  about  four  atmospheres'  pressure  at  about  o°. 
Under  one-atmosphere  pressure  liquid  chlorine  boils  at  —34°. 
This  work  of  Faraday  in  liquefying  chlorine  was  of  very  great 
importance,  since  it  was  the  beginning  of  the  epoch-making 


Halogens  with  Hydrogen  and  Metals 


141 


experiments  in  which  he  succeeded  in  liquefying  all  known 
gases  except  five,  among  which  were  hydrogen,  oxygen,  and 
nitrogen. 

243.  The  Union  of  Chlorine  and  Hydrogen. — Chlorine  and 
hydrogen  do  not  react  at  an  appreciable  rate  at  room  tempera- 
ture if  kept  in  complete  darkness,  but  do  unite  with  explosive 
violence  if  exposed  to  a  bright  light,  hydrogen  chloride  being 

formed,  thus: 

H2+C12->2HC1. 

In  order  to  demonstrate  this  interesting  phenomenon  a  thin- 
walled  glass  bulb  is  filled  with  a  mixture  of  equal  volumes  of 
the  two  gases;  the  bulb  is  then  covered  with  a  thick-walled  bell 
jar  (Fig.  32)  and  strongly  illuminated  either  by  direct  sunlight 


u 

-u- 

II 

•  u 

1 

FIG.  32 


FIG.  33 


or  by  the  rays  from  burning  magnesium  ribbon.  The  sharp 
explosion  which  follows  reduces  the  glass  bulb  to  a  powder,  but 
does  no  damage  to  the  bell  jar.  The  mixture  of  chlorine  and 
hydrogen  is  best  obtained  by  the  electrolysis  of  concentrated 
hydrochloric  acid  in  the  apparatus  shown  in  Fig.  33.  The  inner 
vessel  has  two  carbon  electrodes.  It  is  surrounded  by  a  larger 
vessel,  through  which  water  flows  to  prevent  rise  of  temperature. 
During  the  filling  of  the  bulb  and  up  to  the  time  all  is  ready  for 
the  explosion  it  must  be  shielded  from  bright  light.  The  union  of 
chlorine  with  hydrogen  takes  place  slowly,  without  explosion,  if 
the  mixture  of  the  two  gases  is  exposed  for  a  sufficient  length  of 
time  to  moderate  light  (44). 

244.  The  Burning  of  Hydrogen  in  Chlorine. — If  a  jet  of  hydro- 
gen burning  in  air  is  lowered  into  a  jar  of  chlorine  it  continues 


142 


Introduction  to  General  Chemistry 


r 


FIG.  34 


to  burn  with  a  pale  flame  (Fig.  34).    The  flame  is  the  result  of 

the  intense  heat  produced  by  the  union  of  the  two  gases  to  form 

hydrogen  chloride. 

245.  The  Action  of  Chlorine  on  Water. — Water  dissolves 

about  two  or  three  times  its  own  volume  of  chlorine  at  room 

^ temperature,  giving  a  yellowish  solution 

known  as  chlorine  water.  This  solution 
smells  strongly  of  chlorine  and  is  often 
used  in  the  laboratory  in  place  of 
chlorine  gas.  If  chlorine  water  is 
exposed  to  light  it  soon  loses  its  color 
and  odor,  and  at  the  same  time  a  color- 
less, odorless  gas,  which  proves  to  be 
oxygen,  is  given  off.  The  experiment 
may  readily  be  carried  out  in  the 
manner  shown  in  Fig.  35.  A  cylinder 
filled  with  chlorine  water  is  inverted  in 

a  dish  or  beaker  and  exposed  to  bright  light  for  a  day  or  two. 

The  gas  produced  will  be  found  to  be  oxygen,  formed  according 

to  the  equation 

2C12+2H2O->4HC1+O2. 

This  is  the  reversal  of  the  reaction  by  which  chlorine  is  made  by 
Deacon's  process.  While  chlorine  gas  and  steam  react  only 
partially  at  a  high  temperature,  as  already 
stated,  chlorine  dissolved  in  water  and  ex- 
posed to  light  reacts  slowly,  but  completely,  at 
room  temperature  to  form  hydrochloric  acid 
and  oxygen.  This  curious  difference  in  be- 
havior may  be  traced  to  the  fact  that  while 
gaseous  hydrogen  chloride  and  oxygen  react 
to  the  extent  of  about  80  per  cent  at  400°, 
oxygen  gas  does  not  act  at  all  on  a  solution  of 
hydrochloric  acid  at  room  temperature.  No  chlorine  and  water, 
therefore,  can  be  reproduced  in  cold  water  solution  from  the 
products  of  the  action  of  these  two  substances,  and  so  the 
main  reaction  goes  on  to  completion.  Much  more  is  known 


FIG.  35 


Halogens  with  Hydrogen  and  Metals  143 

about  the  action  of  chlorine  on  water  than  is  contained  in  this 
paragraph,  and  the  subject  will  be  taken  up  again  in  the  following 
chapter. 

246.  The  Union  of  Chlorine  with  Metals.  —  Chlorine  unites 
directly  with  many  metals  forming  chlorides.     In  many  cases 
the  reaction  takes  place  at  once,  with  the  production  of  heat  and 
even  in  some  cases  of  light,  upon  bringing  the  metal  into  chlorine 
gas.     Thin  pieces  of  copper  in  the  form  of  dutch  metal  take  fire 
when  dropped  into  a  jar  of  chlorine,  forming  copper  chloride, 

Cu+Cl2-^CuCl2. 

The  metal  antimony  (symbol  Sb),  in  the  form  of  powder,  also 
unites  with  chlorine,  with  the  production  of  light  and  heat,  if 
sifted  into  a  cylinder  of  the  gas,  antimony  trichloride  being 

formed, 

2Sb+3Cl2->2SbCl3. 

Chlorine  also  unites  directly  with  sodium,  potassium, 
magnesium,  zinc,  iron,  aluminum,  mercury,  and  many  other 
metals  to  form  the  corresponding  chlorides. 

247.  The  Union  of  Chlorine  and  Phosphorus.  —  The  element 
phosphorus  is  a  white,  waxy  solid  which  can  be  made  from  cal- 
cium phosphate,  bone  ash  (158).    We  have  already  seen  (10) 
that  phosphorus  burns  readily  in  the  air.     In  so  doing  it  unites 
with  oxvgen,  thus: 


forming  a  white,  solid  product,  phosphorus  pentoxide.  Phos- 
phorus also  unites  directly  with  chlorine  to  form  either  phos- 
phorus trichloride,  PC13,  or  phosphorus  pentachloride,  PC1S. 
The  preparation  of  the  trichloride  may  be  carried  out  in  a  retort 
as  shown  in  Fig.  36.  About  20  g.  of  dry  phosphorus  are  placed 
in  the  retort  and  a  stream  of  chlorine,  dried  by  passing  it  through 
a  wash  bottle  containing  concentrated  sulfuric  acid,  is  passed 
in  by  means  of  the  glass  tube  which  passes  through  the  stopper 
oi  the  retort.  As  soon  as  the  chlorine  reaches  the  phosphorus, 
union  takes  place  with  the  formation  oj  much  heat  and  the  appear- 
ance of  a  pale  flame.  The  course  of  the  reaction  is  readily  con 
trolled  by  regulating  the  rate  of  flow  of  the  gas  and  by  moving 


144 


Introduction  to  General  Chemistry 


the  gas  inlet  tube  up  or  down  in  the  retort.  If  the  contents  get 
too  hot  so  that  phosphorus  begins  to  distil,  the  temperature  can 
be  lowered  by  raising  the  tube.  On  the  other  hand,  if  yellowish 
crystals  of  the  pentachloride  appear  in  the  retort,  the  tempera- 
ture is  too  low  and  the  tube  should  be  lowered.  The  reaction 
occurs  thus: 

2P  +  3C12->2PC13. 

Phosphorus  trichloride  distils  over  and  condenses  to  a  liquid 
in  the  cooled  receiver.  It  may  be  purified  by  being  distilled 


FIG.  36 

from  a  clean,  dry  retort.  It  is  a  colorless  liquid  which  boils  at 
74°.  It  readily  unites  with  more  chlorine,  forming  solid  crystal- 
line pentachloride,  PC1S: 

PC13+C12->PC1S. 

The  chlorides  of  phosphorus  are  not  salts.     Both  compounds 
are  acted  upon  vigorously  by  water,  according  to  the  following 

equations: 

PC13+3H20->H3P03+3HC1 
PC1S+4H2O->H3P04+5HC1. 

The  products  are  hydrochloric  acid  and  in  the  first  case  phos- 
phorous acid,  H3PO3,  and  in  the  second  case  phosphoric  acid, 
H3P04  (159). 


Halogens  with  Hydrogen  and  Metals  145 

248.  Chlorine   and  Turpentine. — Turpentine   is  a  colorless 
liquid   having   the   formula   d0Hl6.     It   reacts   violently  with 

chlorine,  thus: 

CIOHl6+8Cl2-»ioC+i6HCl. 

The  reaction  is  best  shown  by  bringing  a  strip  of  filter  paper 
which  has  been  dipped  in  turpentine  into  a  cylinder  of  chlorine; 
a  flash  of  flame  occurs  accompanied  by  a  dense,  black  smoke,  due 
to  the  finely  divided  carbon  formed.  This  reaction,  as  well  as 
that  between  chlorine  and  water,  shows  the  great  tendency  of 
chlorine  to  unite  with  hydrogen  even  if  the  hydrogen  is  in  the  form 
of  a  compound. 

249.  Practical  Uses  of  Chlorine. — A  piece  of  litmus  paper 
dipped   into   chlorine  water  becomes   colorless.     Many   other 
vegetable  colors  are  also  bleached  in  the  same  way.     The  process 
is  of  great  practical  importance.     All  white  cotton  goods  have 
been  bleached  by  a  modification  of  this  process,  which  will  be 
described  in  another  chapter  (351). 

In  recent  years  a  new  and  important  use  for  chlorine  has 
been  found  as  a  reagent  for  the  sterilization  of  municipal  water 
supplies.  The  effectiveness  of  chlorine  is  due  to  the  fact  that 
it  is  a  powerful  germicide  by  reason  of  its  great  chemical  activity. 
The  chlorine  is  dissolved  in  the  water  at  the  pumping  stations 
and  during  the  interval  required  for  the  water  to  flow  through 
the  mains  it  reacts  with  the  germs  present  and  is  itself  reduced 
to  harmless  chlorides.  The  water  supply  of  the  city  of  Chicago 
is  purified  in  this  way. 

250.  The    Preparation    of    Hydrochloric    Acid. — We    have 
already  learned  that  hydrogen  chloride  is  made  by  the  action 
of  sulfuric  acid  on  common  salt.    The  best  laboratory  method 
is  that  described  earlier  (103),  the  reaction  taking  place  according 
to  the  following  equation: 

NaCl+H2SO4->NaHSO4+HCl. 

If,  however,  double  the  proportion  of  salt  indicated  by  this  equa- 
tion is  taken  and  the  temperature  is  finally  raised  sufficiently, 
the  following  reaction  will  take  place: 

2NaCl-f  H2S04->Na2SO4+  2HC1. 


146  Introduction  to  General  Chemistry 

By  the  last  reaction  a  given  quantity  of  sulfuric  acid  will  produce 
double  the  quantity  of  hydrogen  chloride  as  in  the  first;  it  is 
therefore  the  more  economical  and  is  the  one  used  in  the  com- 
mercial production  of  hydrochloric  acid. 

The  union  of  hydrogen  and  chlorine  to  form  hydrogen 
chloride  has  already  been  discussed  (44,  243).  In  recent  years, 
since  chlorine  has  become  available  in  immense  quantities  as 
a  by-product  of  the  manufacture  of  caustic  soda,  some  hydro- 
chloric acid  has  been  produced  commercially  in  this  way. 

The  old  name  for  hydrochloric  acid  was  muriatic  acid,  and 
this  is  the  name  by  which  the  crude  acid  is  still  commonly  known 
in  trade. 

251.  The  Physical  Properties  of  Hydrogen  Chloride.— 
Hydrogen  chloride  is  a  colorless  gas,  having  a  choking  odor  and 
forming  a  cloud  of  white  fumes  in  moist  air.  Its  density  is 
somewhat  greater  than  that  of  air;  one  liter  weighs  i.642g. 
The  gas  is  very  soluble  in  water;  at  room  temperature  water 
dissolves  about  450  times  its  volume  of  the  gas,  giving  a  con- 
centrated solution  of  hydrochloric  acid.  Considerable  heat  is 
produced  when  the  gas  dissolves  in  water,  so  that  the  solution 
becomes  decidedly  warm.  In  general,  when  gases  dissolve  in 
water  heat  is  produced.  So-called  chemically  pure  hydrochloric 
acid  has  a  specific  gravity  of  i .  2  and  contains  about  37  per  cent 
of  hydrogen  chloride,  the  balance  being  water. 

When  the  37  per  cent  acid  is  heated,  hydrogen  chloride 
gas  is  given  off,  together  with  some  water  vapor,  and  the  remain- 
ing  solution  becomes  less  concentrated.  Finally  the  tempera- 
ture rises  to  110°  before  the  liquid  boils;  by  this  time  the 
concentration  has  decreased  to  20  per  cent.  As  the  solution 
continues  to  boil,  its  concentration,  20  per  cent,  and  boiling- 
point,  110°,  remain  constant;  the  condensed  vapor,  the  so-called 
distillate)  also  has  a  concentration  of  20  per  cent. 

On  the  other  hand,  if  very  dilute  hydrochloric  acid  is  boiled 
it  loses  water  chiefly  and  becomes  more  concentrated;  finally, 
when  the  concentration  has  reached  20  per  cent  the  boiling 
temperature  has  become  110°,  after  which  both  concentration 
and  boiling-point  remain  constant. 


Halogens  with  Hydrogen  and  Metals  147 

252.  The  Chemical  Properties  of  Hydrochloric  Acid. — The 

most  important  chemical  properties  of  hydrochloric  acid  have 
already  been  studied.  These  may  be  briefly  reviewed  in  this 
paragraph.  Hydrochloric  acid  is  perhaps  the  most  typical  of 
all  acids;  it  turns  litmus  red  and  its  very  dilute  solution,  say 
i  per  cent,  has  a  pleasant  sour  taste ;  it  neutralizes  the  hydroxides 
and  oxides  of  metals,  forming  chlorides  and  water,  for  example: 

NaOH+HCl->NaCl+H2O 
CuO+2HCl->CuCl2+H2O. 

It  acts  on  many  metals  forming  chlorides  and  hydrogen,  thus : 
Fe+2HCl->FeCl2+H2. 

The  addition  of  hydrochloric  acid  to  solutions  of  salts  of 
silver  (169),  lead  (167),  and  univalent  mercury  (182)  gives 
precipitates  of  insoluble  chlorides,  thus: 

AgNO3+HCl->AgCl+HNO3. 

Oxidizing  agents,  such  as  oxygen  gas  at  a  high  temperature 
and  higher  oxides  of  the  metals  like  manganese  dioxide,  liberate 

chlorine : 

4HC1+O2->2C12+  2H2O  (239) 

4HCl+MnO2->Cl2+MnCl2+2H2O.  (235) 

Hydrochloric  acid  is  an  almost  indispensable  chemical 
reagent.  It  is  used  extensively  both  in  scientific  and  in  technical 
work.  It  is  manufactured  in  large  quantities  and  is  an  impor- 
tant article  of  commerce. 

253.  The  Action  of  Hydrochloric  Acid  on  Sodium  Hydrogen 
Sulfate. — If  concentrated  hydrochloric  acid  is  added  slowly, 
with  stirring,  to  a  concentrated  solution  of  sodium  hydrogen 
sulfate,  a  white  crystalline  precipitate  is  formed,  which,  when 
filtered  out,  washed  with  a  little  water,  and  dried,  is  found  to 
consist  of  pure  sodium  chloride.     The  reaction  is  represented 

thus: 

HCl+NaHSO4->NaCl+H2SO4. 

This  is  seen  to  be  just  the  reverse  of  the  reaction  by  which 
hydrogen  chloride  is  made  from  salt.  It  is  therefore  a  reversible 


148  Introduction  to  General  Chemistry 

reaction.  The  direction  which  the  reaction  will  take  depends 
upon  the  amount  of  water  present  and  the  temperature.  Dry 
salt  and  anhydrous  (water-free)  sulfuric  acid  react  practically 
completely  to  form  hydrogen  chloride  and  sodium  hydrogen 
sulfate;  while  sufficiently  dilute  sulfuric  acid  and  salt  do  not 
give  off  any  hydrogen  chloride  gas.  The  reason  is  simple:  the 
gas  is  very  soluble  in  water,  and  even  if  it  were  formed  it  would 
remain  dissolved  in  the  water  present.  The  fact  that  con- 
centrated solutions  of  hydrogen  chloride  and  sodium  hydrogen 
sulfate  give  a  precipitate  qf  solid  sodium  chloride  shows  clearly 
that  the  reaction  has  a  tendency  to  reverse.  It  seems  probable 
that  in  the  presence  of  much  water,  that  is,  in  dilute  solution, 
all  four  of  the  substances  are  present  in  any  solution  that  is  made 
by  bringing  either  pair  of  substances  together.  In  such  a  solution 
we  may  say  that  there  exists  a  state  of  equilibrium  as  the  result 
of  each  pair  of  substances  on  the  same  side  of  the  equation  con- 
tinuously reacting  to  form  the  pair  on  the  opposite  side,  thus: 

H2S04+NaCl^NaHSO4+HCl. 

254.  Bromine. — The  element  bromine  (symbol  Br)  resembles 
chlorine  more  closely  than  does  any  other  element.     It  does  not 
%occur  free  in  nature.     Its  salts,  the  bromides,  are  frequently 
found  in  small  amounts  associated  with  chlorides.     Sea-water 
contains  a  small  proportion  of  bromides.    Large  quantities  of 
bromides  are  obtained  from  deposits  accompanying  those  of 
sodium  nitrate  in  the  desert  regions  of  Chile.     The  brines  from 
salt  springs  in  Michigan  also  furnish  bromides  in  commercial 
quantities. 

255.  Sodium  bromide,  NaBr,  potassium  bromide,  KBr,  and 
magnesium  bromide,  MgBr2,  are  the  commonest  salts  directly 
obtainable  from  natural  salt  deposits  and  brines.     From  any 
of  these  the  element  is  readily  set  free  by  the  action  of  chlorine, 

thus: 

2KBr+Cl2->2KCl+Br2. 

Upon  passing  chlorine  gas  into  a  solution  of  potassium  bromide, 
the  solution  turns  brown  and  when  heated  gives  off  reddish-brown 
vapors  of  bromine,  which  when  cooled  condense  to  liquid 


Halogens  with  Hydrogen  and  Metals 


149 


bromine.  Bromine  is  a  reddish-brown  liquid  which  has  a 
density  over  three  times  that  of  water.  It  boils  at  58°  and  readily 
volatilizes  at  ordinary  temperatures.  The  vapor  is,  if  anything, 
more  irritating  to  mucous  membranes  than  chlorine,  and  the 
liquid  produces  deep  burns  when  brought  into  contact  with  the 
skin.  Bromine  must  be  handled  with  extreme  caution.  In  case 
of  accident  wash  off  the  bromine  with  water  immediately;  then 
consult  an  instructor  regarding  further  treatment. 

Bromine  dissolves  in  water  to  the  extent  of  about  3  per  cent 
to  form  a  light-brown  solution,  known  as  bromine  water. 

256.  Hydrobromic  Acid,  HBr.  —  Hydrogen  bromide,  the 
water  solution  of  which  is  known  as  hydrobromic  acid,  can  be 
made  by  the  direct  union  of  its  constituent  elements: 

H2+Br2->2HBr. 

The  best  method  of  making  hydrogen  bromide  is  based  on  the  fact 
that  bromine  unites  with  phosphorus  to  form  a  tribromide  or  a 
pentabromide,  thus, 


These  compounds  are  entirely  analogous  to  PC13  and  PC1S  (247). 
The  bromides  of  phosphorus  also  resemble  the  chlorides  in  their 
reactions  with  water,  thus  : 

PBr3+3H2O^H3P03+3HBr 
PBrs-f-4H2O->H3PO4 


The  preparation  of  hydrobromic  acid  is  carried  out  in  the 
apparatus  shown  in  Fig.  37. 

Ten  grams  of  red  phosphorus, 
10  c.c.  of  water,  and  20  to  25  g. 
of  quartz  sand  are  placed  in  a 
250  c.c.  flask  and  15  c.c.  of  bro- 
mine, contained  in  the  dropping 
funnel,  are  allowed  to  run  in 
slowly,  drop  by  drep.  The  U-tube 
contains  some  pieces  of  broken  FIG 

glass    or    brick    or    similar  inert 
material  mixed  with  3  or  4  g.  of  red  phosphorus,  the  object  of 


150  Introduction  to  General  Chemistry 

the  glass  or  brick  being  to  distribute  the  phosphorus  so  that  it 
will  present  the  maximum  of  surface.  The  hydrogen  bromide 
given  off  .is  freed  from  accompanying  bromine  vapor  by  the 
phosphorus  in  the  U-tube  and  is  absorbed  by  water  contained 
in  the  cylinder.  The  delivery  tube  should  not  dip  into  the 
water  in  the  cylinder,  since  the  gas  is  so  soluble  that  there  would 
be  danger  of  water  getting  back  into  the  U-tube  and  flask. 

257.  The  Properties  of  Hydrogen  Bromide.  —  Hydrogen 
bromide  is  a  colorless  gas  with  a  choking  odor;  it  gives  white 
fumes  in  moist  air  and  dissolves  abundantly  in  water  to  form  a 
solution  known  as  hydrobromic  acid.  This  is  a  colorless  liquid 
which  closely  resembles  hydrochloric  acid  in  its  properties.  It 
neutralizes  bases  and  unites  with  metallic  oxides  to  form  salts 
called  bromides,  for  example  : 

NaOH+HBr->NaBr+H2O 


CuO+  2HBr^CuBr2+H2O 
Al(OH)3+3HBr->AlBr3+3H20. 

The  bromides  of  silver,  lead,  and  univalent  mercury  are  almost 
insoluble  in  water,  as  are  the  chlorides  of  these  same  metals  (252). 
All  other  bromides  are  easily  soluble.  The  addition  of  hydro- 
bromic acid  or  any  soluble  bromide  to  a  solution  of  a  salt  of 
silver,  lead,  or  univalent  mercury  gives  a  white  precipitate  of 
the  insoluble  bromide,  thus: 

Pb(NO3)2+  2HBr->PbBr2+  2HNO3. 

258.  The  Oxidation  of  Hydrobromic  Acid.  —  Hydrogen  bro- 
mide and  oxygen  gases  react  when  heated  to  form  bromine  and 
water, 


This  reaction  is  analogous  to  that  between  hydrogen  chloride 
and  oxygen  (239),  but  takes  place  far  more  completely,  indi- 
cating that  hydrogen  bromide  is  more  easily  oxidized  than  hy- 
drogen chloride.  Other  oxidizing  agents,  such  as  manganese 
dioxide,  readily  set  free  bromine: 

4HBr+Mn02-»MnBr2+Br2+  2H2O. 


Halogens  with  Hydrogen  and  Metals  151 

In  the  technical  preparation  of  bromine  by  means  of  this 
reaction  sodium  bromide  is  treated  with  dilute  sulfuric  acid 
and  manganese  dioxide.  In  this  case  all  of  the  available 
bromine  is  set  free. 

2NaBr+2H2SO4+MnO2->Na2SO4+MnSO4+Br2+2H2O. 

259.  The  Action  of  Chlorine  on  Bromides. — A  solution  of 
any  bromide  reacts  with  chlorine  to  form  a  chloride  and  free 

bromine, 

2KBr+Cl2->2KCl+Br2. 

Similarly,  hydrobromic  acid  and  chlorine  give  hydrochloric  acid 
and  bromine.  These  reactions  are  nearly  complete,  that  is, 
they  are  not  reversible  to  any  marked  extent,  so  that  we  may 
conclude  that  the  metals  and  hydrogen  form  by  preference  com- 
pounds with  chlorine  rather  than  with  bromine.  This  fact  may 
also  be  expressed  by  saying  that  chlorine  has  greater  affinity  than 
bromine  for  metals  and  hydrogen.  Using  this  mode  of  expres- 
sion, we  should  also  say  that  oxygen  has  greater  affinity  than 
bromine  for  hydrogen,  since  hydrogen  bromide  and  oxygen  give 
water  and  free  bromine. 

260.  The  Uses  of  Bromine  and  Its  Compounds. — Potassium 
and  sodium  bromides  are  used  extensively  in  medicine  as  seda- 
tives.    Silver  bromide  is  the  light-sensitive  substance  of  photo- 
graphic plates.    The  free  element  is  extensively  used  in  the 
manufacture  of  important  coal-tar  dyes. 

261.  Iodine. — The  element  iodine  (symbol  I),  bears  almost 
the  same  relation  to  bromine  that  the  latter  bears  to  chlorine. 
It  does  not  occur  free  in  nature,  but  is  readily  prepared  from 
its  compounds,  the  iodides  of  sodium  or  potassium,  which  are 
obtained  from  two  principal  natural  sources. 

Certain  seaweeds  contain  small  amounts  of  combined  iodine 
which  has  been  taken  up  from  sea-water  in  which  a  minute 
quantity  is  present.  The  ashes  left  upon  burning  the  dried 
seaweed  yield  by  extraction  with  water  sodium  iodide,  Nal, 
and  potassium  iodide,  KI.  Iodine  compounds  are  also  obtained 
as  by-products  in  the  purification  of  the  sodium  nitrate  found 


152  Introduction  to  General  Chemistry 

in  Chile  (104).     Iodine  is  set  free  from  iodides  by  the  action  of 

chlorine,  thus: 

2NaI+Cl2->2NaCl+I2. 

It  is  also  liberated  by  the  action  of  manganese  dioxide  and 
sulfuric  acid. 

262.  The  Physical  Properties  of  Iodine.  —  Iodine  is  an  almost 
black,  crystalline  substance,  having  a  density  of  nearly  five.     It 
melts  at  114°  and  boils  at  a  somewhat  higher  temperature,  pro- 
ducing a  vapor  having  a  magnificent  violet  color.    At  a  tempera- 
ture slightly  below  its  melting-point  iodine  has  so  great  a  vapor 
pressure  that  by  cautious  heating  it  may  be  volatilized  com- 
pletely without  being  melted.     If  the  vapor  is  allowed  to  strike 
a  cold  surface  crystals  of  iodine  deposit  directly  without  pre- 
liminary formation  of  liquid  iodine.     The  sublimation  (179)  of 
iodine  in  this  way  is  an  important  step  in  the  purification  of  this 
element. 

Iodine  is  very  slightly  soluble  in  water,  giving  a  faintly 
brownish  solution.  It  dissolves  abundantly  in  water  solutions 
of  potassium  or  sodium  iodide.  It  dissolves  easily  in  alcohol, 
forming  a  dark-brown  solution  called  by  druggists  tincture  of 
iodine.  Iodine  also  dissolves  easily  in  ether,  forming  a  brown 
solution,  and  in  chloroform  and  carbon  disulfide,  forming 
violet-colored  solutions. 

263.  Iodine  and  Starch.  —  If  a  dilute  solution  of  iodine  is 
added  to  water  containing  a  little  starch  paste,  made  by  boiling 
starch  with  50  to  100  times  its  weight  of  water,  a  deep  blue- 
colored  solution  results.     This  reaction  is  a  characteristic  and 
very  delicate  test  for  free  iodine.     Iodides,  like  KI,  do  not  give 
this  test;  but  by  adding  chlorine  to  a  solution  of  an  iodide  the 
element  is  set  free  and  can  then  be  recognized  by  the  starch  test. 
An  excess  of  chlorine  interferes  with  this  test. 

264.  Hydrogen   Iodide,   HI.  —  Iodine   and    hydrogen   unite 
slowly  at  a  temperature  of  400°  to  form  hydrogen  iodide,  thus: 


The  product  is  a  colorless  gas,  analogous  to  hydrogen  chloride 
and  hydrogen  bromide.  Like  these  latter  gases  it  dissolves 
abundantly  in  water,  and  forms  fumes  in  moist  air. 


Halogens  with  Hydrogen  and  Metals  153 

Hydrogen  iodide  is  easily  made  by  a  reaction  resembling 
that  used  for  making  hydrogen  bromide.  Iodine  forms  with 
phosphorus  a  tri-iodide,  PI3.  This  reacts  with  water  to  form 
phosphorous  and  hydriodic  acids  thus: 

PI3+3H20->H3P03+3HI. 

The  process  of  making  hydrogen  iodide  is  carried  out  by  placing 
a  mixture  of  powdered  iodine  and  red  phosphorus  in  a  flask  and 
running  in  water,  drop  by  drop  from  a  dropping  funnel,  care 
being  taken  not  to  use  more  water  than  is  necessary,  since  an 
excess  of  water  would  dissolve  the  gas  and  so  prevent  its  escape 
from  the  flask.  The  apparatus  used  for  making  hydrogen 
bromide,  Fig.  37,  may  be  used  in  this  case.  The  U-tube  con- 
taining red  phosphorus  serves  here  to  remove  iodine  vapor. 
The  hydrogen  iodide  gas  may  be  collected  by  downward  dis- 
placement of  air  or  it  may  be  dissolved  in  water  to  form  a  solu- 
tion of  hydriodic  acid. 

265.  Hydriodic  acid  is  colorless  when  pure,  but  is  brown  if 
it  contains  free  iodine,  which  it  dissolves  readily.  It  neutralizes 
bases  and  so  yields  salts  called  iodides,  for  example: 

HI+NaOH-»NaI-f-H2O 
2HI+Ca(OH)2->CaI2-f2H2O. 

Hydriodic  acid  acts  on  metals  similarly  to  hydrochloric  acid, 
giving  iodides  and  hydrogen,  thus: 


Hydriodic  acid  is  much  more  easily  oxidized  than  is  hydro- 
bromic  acid,  which  in  turn  is  more  easily  oxidized  than  hydro- 
chloric acid;  while  all  three  acids  are  oxidized  by  powerful 
oxidizing  agents  such  as  manganese  dioxide  and  lead  dioxide; 
hydriodic  acid,  even  in  dilute  solution,  is  oxidized  slowly  by 
atmospheric  oxygen,  which  has  no  action  whatever  on  dilute 

hydrochloric  acid: 

4HI+O2->2H2O+2l2. 

The  iodine  which  is  slowly  liberated  according  to  the  equation 
given  above  remains  dissolved  in  the  unchanged  acid  and  gives 
it  a  brown  color. 


154  Introduction  to  General  Chemistry 

266.  Uses  of  Iodine  and  Iodides.  —  Iodine  is  used  extensively 
in  certain  processes  of  analysis  and  also  in  the  preparation  of 
important  compounds  containing  the  element  carbon,  so-called 
organic  compounds.     Iodine  in  the  form  of  tincture  of  iodine, 
which  is  a  solution  of  iodine  in  alcohol,  is  used  externally  as  an 
antiseptic  and  also  as  a  counterirritant  in  medicine.    The  iodides 
of  potassium,  sodium,  and  ammonium  are  of  great  importance 
for  internal  administration  in  medicine. 

267.  Fluorine.  —  The  element  fluorine  (symbol  F),  is  classed 
among  the  halogens,  although  it  is  less  closely  related  to  the 
other  three  halogens,  chlorine,  bromine,  and  iodine,  than  these 
three  are  to  one  another.    The  atomic  weights  of  these  elements 
are:    fluorine,   19;    chlorine,  35.5;    bromine,  80;    iodine,  127. 
Fluorine  has,  therefore,  the  smallest  atomic  weight  of  any  of 
the  halogens.     We  might  expect  it  to  resemble  chlorine  more 
closely  than  it  does  bromine  and  iodine  and,  in  fact,  such  is  the 
case.     It  is  a  pale-yellow  gas  which  is  very  active  chemically  and 
never  occurs  free  in  nature.     Its  most  abundant  natural  com- 
pound is  calcium  fluoride  or  fluor-spar,  CaF2.    It  also  occurs  as 
cryolite,  sodium   aluminum  fluoride,  3NaF'AlF3.    These  sub- 
stances are  salts  of  hydrofluoric  acid.    We  might  expect  that 
free  fluorine  could  be  made  by  oxidizing  hydrofluoric  acid  with 
manganese  dioxide,  thus: 

4HF+MnO2->MnF2+F2+2H2O; 

but  we  find,  in  fact,  that  hydrofluoric  acid  is  entirely  unacted 
upon  by  the  most  powerful  oxidizing  agents.  The  free  element 
was  first  made  by  Moissan,  by  the  electrolysis  of  anhydrous 
liquid  hydrogen  fluoride,  in  which  some  potassium  fluoride,  KF, 
was  dissolved  to  make  it  conduct  electricity  readily.  The  pro- 
ducts of  the  electrolysis  were  fluorine  and  hydrogen  : 


Fluorine  is  one  of  the  most  active  of  all  elements.  It  rapidly 
attacks  glass  and  also  most  metals,  and  it  reacts  at  once  with 
water  forming  hydrofluoric  acid  and  oxygen: 

2F2+2H2O->4HF+O2. 


Halogens  with  Hydrogen  and  Metals  155 

The  preparation  of  fluorine  is  a  matter  of  great  difficulty,  for 
which  reason  it  is  very  seldom  made. 

268.  Hydrogen   Fluoride,   HF. — Hydrogen  fluoride,   a  gas 
whose  water  solution  is  called  hydrofluoric  acid,  is  the  most 
important  compound  of  fluorine.     It  is  formed  by  the  action  of 
concentrated  sulfuric  acid  on  powdered  calcium  fluoride: 

H2SO4+  CaF2-»CaSO4+  2HF. 

It  is  a  colorless  gas  with  a  choking  odor.  At  temperatures  of 
100°  and  higher  its  density  shows  that  the  gas  has  the  formula 
HF;  at  room  temperature  the  density  is  more  than  double 
that  expected  for  a  gas-  with  the  formula  HF.  This  fact  leads 
to  the  conclusion  that  the  single  molecules  have  become  asso- 
ciated, probably  to  form  double  or  triple  molecules  such  as 
H2F2  and  H3F3.  Hydrogen  fluoride  gas  is  condensed  to  a  liquid 
merely  by  cooling  it  with  ice ;  colorless  liquid  hydrogen  fluoride, 
so  obtained,  boils  at  19°. 

269.  Hydrofluoric  Acid  and  Its  Salts. — A  30  per  cent  solu- 
tion of  hydrofluoric  acid  is  an  important  article  of  commerce. 
The  acid  has  several  practical  uses.     These  include  the  etching 
and  polishing  of  glass,  the  removal  of  sand  from  castings,  and 
the  preparation  of  its  salts  and  also  of  hydrofluosilicic  acid, 
H2SiF6. 

Hydrofluoric  acid  forms  with  bases  salts  called  fluorides. 
The  soluble  fluorides  are  very  effective  preservatives,  since  they 
inhibit  the  growth  of  bacteria,  molds,  etc.  But  their  use  in 
foodstuffs  is  prohibited  because  of  their  interference  with  diges- 
tion. 

Ammonium  fluoride,  NH4F,  is  used  as  a  disinfectant  for 
utensils  used  in  breweries.  Sodium  fluoride,  NaF,  is  extensively 
used  as  a  vermin  exterminator  for  poultry. 

270.  The  Action  of  Hydrogen  Fluoride  on  Quartz. — We  must 
now  digress  a  little  from  the  subject  in  hand  in  order  to  be  able 
fully  to  understand  one  of  the  most  interesting  reactions  of 
hydrogen  fluoride.     The  substance  called  quartz  is  the  oxide 
of  an  element  silicon  (symbol  Si)  and  has  the  formula  Si02. 
Common  sand  is  more  or  less  pure  quartz.     Glass,  which  is 


156  Introduction  to  General  Chemistry 

made  by  melting  together  sand,  sodium  carbonate,  and  slaked 
lime,  may  be  considered  a  mixture  of  sodium  silicate,  Na2Si03, 
and  calcium  silicate,  CaSiO3.  Hydrofluoric  acid  and  quartz 
react  very  readily  to  form  gaseous  silicon  fluoride,  SiF4,  and 

water,  thus: 

4HF+SiO2->SiF4+2H20. 

This  is  a  very  characteristic  reaction;  none  of  the  other  halogen 
acids  have  any  action  on  quartz. 

Glass,  which  is  almost  unaffected  by  the  other  halogen  acids, 
is  rapidly  attacked  by  either  hydrogen  fluoride  gas  or  hydro- 
fluoric acid  solution.  The  fluorine  unites,  not  only  with  the 
silicon,  as  in  the  case  of  quartz,  forming  silicon  fluoride,  but  also 
with  the  sodium  and  calcium  forming  sodium  fluoride,  NaF,  and 
calcium  fluoride,  CaF2,  the  reactions  being: 

Na2SiO3+6HF->SiF4+2NaF+3H20 
CaSiO3+6HF->SiF4+CaF2+3H2O. 

The  result  is  that  glass  dissolves  very  easily  in  hydrofluoric 
acid,  in  consequence  of  which  this  acid  cannot  be  kept  in  glass 
bottles.  Parafline  and  other  waxes,  which  are  not  attacked,  are 
used  for  bottles  for  this  acid,  while  larger  containing  vessels  are 
made  of  lead. 

271.  Etching  Glass  with  Hydrogen  Fluoride. — The  etching  of 
glass  may  be  illustrated  by  coating  a  glass  plate  with  a  thin 
layer  of  paraffine,  and  after  making  a  design  or  inscription  by 
means  of  a  hard  pencil  which  will  cut  through  the  paraffine  and 
thus  expose  the  surface  of  the  glass,  exposing  the  plate  to  the 
action  of  hydrogen  fluoride  gas.    The  gas  is  easily  made  by 
mixing  a  few  grams  of  powdered  fluor  spar  with  concentrated 
sulfuric  acid  in  a  shallow  lead  dish.     The  latter  is  covered  with 
the  glass  plate  and  set  aside  for  ten  or  fifteen  minutes.     Upon 
removing  the  paraffine,  the  design  will  be  found  to  have  been 
etched  upon  the  glass. 

272.  Hydrofluosilicic  Acid,  H2SiF6. — Hydrogen  fluoride  and 
silicon  tetrafluoride  unite  readily  in  the  presence  of  water  to 
form  a  solution  of  hydrofluosilicic  acid: 

2HF-j-SiF4-»H2SiF6. 


Halogens  with  Hydrogen  and  Metals  157 

The  solution  is  a  colorless,  odorless  liquid  which  does  not  attack 
glass  appreciably.  It  has  well-characterized  acid  properties: 
it  reddens  litmus,  has  a  sour  taste,  and  neutralizes  bases  to  form 
salts.  This  acid  is  important  technically.  It  is  made,  in  practice, 
by  the  action  of  hydrofluoric  acid  solution  on  quartz  sand : 

6HF+  SiO2->H2SiF6+  2H2O. 

The  acid  is  used  for  the  preparation  of  its  sodium,  magnesium, 
and  lead  salts.  Sodium  fluosilicate,  Na2SiF6,  is  extensively  used 
in  making  white  enameled  ware  and  also  white,  or  so-called  milk, 
glass.  It  is  remarkable  in  being  one  of  the  very  few  nearly 
insoluble  salts  of  sodium.  It  is  obtained  as  a  white  precipitate 
when  solutions  of  common  salt  and  hydrofluosilicic  acid  are  mixed. 

2NaCl+H2SiF6->Na2SiF6+  2HC1. 

Magnesium  fluosilicate,  MgSiF6,  easily  soluble  in  water,  is 
used  to  harden  concrete.  Lead  fluosilicate,  PbSiF6,  also  easily 
soluble  in  water,  is  made  as  an  intermediate  product  in  refining 
lead  (Betts's  process). 


CHAPTER  XIII 
CHEMICAL  EQUILIBRIUM 

273.  Incomplete  Physical  Processes. — While  many  physical 
processes  are  seemingly  complete,  there  are  others  which  stop 
far  short  of  completion.     Thus,  for  example,  if  a  small  bulb  of 
water  is  broken  in  a  large  closed  bottle,  evaporation  of  the 
water  will  start  at  once,  but  will  apparently  cease  as  soon  as  the 
pressure  of  the  vapor  reaches  a  value  which  is  definite  for  a 
definite  temperature,  although  much  liquid  may  still  remain 
(112). 

If  we  add  to  some  water  an  equal  weight  of  common  salt, 
the  latter  will  at  once  start  to  dissolve  and  will  continue  to  do  so 
until  the  solution  has,  for  a  given  temperature,  a  certain  definite 
concentration;  then,  although  much  solid  salt  is  still  present, 
no  further  increase  in  concentration  will  take  place  (122). 

When  water  in  a  closed  vessel,  which  it  fills  but  partially, 
reaches  its  maximum  vapor  pressure  for  a  given  temperature, 
we  believe  (201)  that  for  every  molecule  that  passes  from  liquid 
to  vapor  there  is  one  that  passes  from  vapor  to  liquid.  We  say 
that  there  is  equilibrium  between  liquid  and  vapor.  We  believe 
that  a  similar  condition  exists  when  a  solid  apparently  stops 
dissolving  in  a  solution  (207).  The  apparent  state  of  rest  or 
inaction  in  both  cases  is  very  probably  one  in  which  two  opposing 
actions  exactly  counteract  the  effects  of  each  other. 

274.  Incomplete  Chemical  Reactions. — Just  as  in  the  case 
of  physical  processes,  there  are  also  some  chemical  reactions 
that  do  not  go  to  completion.     We  have  already  studied  some 
reactions  of  this  kind  and  must  now  consider  the  matter  more 
fully,  as  it  is  one  of  great  importance. 

The  reaction  between  hydrogen  chloride  and  oxygen  at  400° 
has  been  considered  (239,  240)  under  the  heading  "  Deacon's 
Process."  It  has  been  pointed  out  (245)  that  only  80  per  cent 
of  the  hydrogen  chloride  is  oxidized  when  a  mixture  of  this  gas 

158 


Chemical  Equilibrium  159 

is  heated  with  oxygen  in  the  proportion  indicated  in  the  following 
equation  : 


On  the  other  hand,  when  a  mixture  of  two  formula  weights 
each  of  chlorine  and  water  is  also  heated  to  400°,  80  per  cent  of 
the  chlorine  remains  unchanged,  while  20  per  cent  is  converted 
into  hydrogen  chloride.  It  thus  happens  that  whether  we  start 
with  the  pair  of  gases  on  the  left  side  of  the  foregoing  equation 
or  the  pair  on  the  right,  taking  in  each  case  the  amounts  indi- 
cated in  this  equation,  there  results  a  mixture  of  the  four  gases 
which  has  exactly  the  same  amount  of  each  gas  present  in  the 
two  cases.  It  is  easy  to  see  that  the  cause  of  each  reaction  being 
incomplete  is  found  in  the  fact  that  the  products  of  either  reac- 
tion again  react  in  the  opposite  direction.  In  the  mixture  of  the 
four  gases  which  finally  results  we  say  that  a  state  of  equilibrium 
exists  and  that  the  apparent  stopping  of  further  change  is  really 
the  result  of  the  formation  of  hydrogen  chloride  and  oxygen  at 
just  the  same  rate  as  that  at  which  these  two  gases  change  into 
chlorine  and  water. 

275.  Velocity  of  Chemical  Change.  —  The  idea  that  a  state 
of  chemical  equilibrium  is  the  result  of  two  opposing  changes 
which  take  place  continuously  at  such  rates  or  with  such  velocities 
that  for  every  molecule  of  a  given  substance  formed  one  also 
disappears  would  imply  that  chemical  changes  take  place  gradually 
and  possibly  at  definite  speeds  or  velocities. 

It  is  well  known  that  certain  reactions,  as  for  example  the 
burning  of  a  candle  or  the  action  of  an  acid  on  a  metal,  certainly 
do  take  place  gradually.  It  is  not  so  plain  that  if  the  reaction 
takes  place  between  two  perfectly  mixed  gases  or  between  two 
substances  completely  dissolved  as  a  uniformly  mixed  solution 
that  time  is  required  for  the  reaction  to  take  place.  Nevertheless 
it  is  probable  that  no  reaction,  even  an  explosion,  however  rapid 
it  may  be,  is  absolutely  instantaneous. 

The  speed  or  velocity  of  reaction  in  a  uniformly  mixed  solu- 
tion may  be  beautifully  and  convincingly  demonstrated  by  means 
of  the  following  experiment: 


Introduction  to  General  Chemistry 

To  800  c.c.  of  water  contained  in  a  flask  there  is  added  25  c.c. 
of  starch  solution  (made  by  boiling  2  g.  of  starch  with  100  c.c. 
of  water)  and  15  c.c.  of  a  3  per  cent  solution  of  iodic  acid,  HIO3. 
The  solution  is  then  well  mixed  and  5  c.c.  of  a  3  per  cent  solution 
of  sulfur  dioxide,  S02,  is  added  and  the  contents  of  the  flask 
are  at  once  thoroughly  mixed  by  being  shaken.  The  time  of 
adding  the  sulfur  dioxide  solution  is  accurately  noted — best  with 
a  stop  watch.  No  change  will  be  seen  in  the  colorless  solution 
for  about  60  seconds,  then  the  whole  solution  will  suddenly  turn 
deep  blue.  The  result  is  startling! 

If  the  experiment  is  repeated,  using  the  same  amounts  of 
water  and  of  each  of  the  three  solutions,  and  if  the  temperature 
is  also  the  same,  it  will  be  found  that  the  time  required  for  the 
change  to  occur  is  always  the  same.  If,  however,  we  increase 
the  amount  of  sulfur  dioxide  solution  added  from  5  c.c.  to 
10  c.c. ,  everything  else  remaining  the  same,  the  time  required 
for  the  change  will  be  decreased  to  about  30  seconds.  The 
increased  "velocity  is  the  result  of  the  increase  in  concentration  of 
the  sulfur  dioxide. 

276.  The  Effect  of  Temperature  on  Reaction  Velocity.— The 
effect  on  the  velocity  of  increasing  the  temperature  is  easily 
shown  by  starting  with  water  at  25°  instead  of  at  20°,  when  it 
will  be  clear  that  at  the  higher  temperature  the  velocity  is  decidedly 
greater. 

277.  The  Action  of   Sulfur   Dioxide  on  Iodic  Acid. — The 
chemical  changes  involved  in  the  reaction  just  described  need 
not  greatly  concern  the  student  at  this  time,  as  they  are  of  less 
importance  than  the  main  facts  of  reaction  velocity  that  they 
serve  here  to  illustrate.     But  as  it  is  only  natural  to  wonder 
what  has  happened  in  such  a  striking  experiment,  the  equations 
for  the  reactions  may  now  be  given.     In  the  first  place,  sulfur 
dioxide,  S02,  and  water  form  sulfurous  acid,  H2SO3, 

SO2+H2O->H2SO3. 

The  latter  reacts  with  the  iodic  acid,  forming  hydriodic  and 
sulfuric  acids,  thus: 

HI03-f-3H2S03-»HI+3H2S04. 


Chemical  Equilibrium  161 

But  hydriodic  acid  can  also  react  with  iodic  acid  to  form  free 
iodine  and  water, 

HI03+5HI->3I2+3H20, 

and  then  the  iodine  set  free  acts  on  the  starch  to  produce  the 
blue  color.  Now  this  reaction  between  iodic  and  hydriodic 
acid  does  not  take  place  until  all  the  sulfurous  acid  has  dis- 
appeared. The  time  observed  for  the  appearance  of  the  blue 
color  is  therefore  essentially  that  required  for  the  complete 
oxidation  of  the  sulfurous  acid. 

278.  The  Kinetic  Hypothesis  Applied  to  Reaction  Velocity.— 
The  application  of  the  kinetic-molecular  hypothesis  (chap,  x) 
leads  to  a  simple  and  reasonable  explanation  of  reaction  velocity. 

Let  us  suppose  that  two  gases,  A  and  B,  can  unite  to  form 
a  compound  AB,  and  let  the  reaction  be  represented  by  the 
equation 

A+B+AB. 

Let  us  also  suppose  that  the  reaction  takes  place  rather 
slowly  after  the  two  gases  have  been  thoroughly  mixed.  We 
may  now  consider  what  determines  the  rate  at  which  A  and  B 
unite.  It  is  obvious  that  union  can  occur  only  when  a  molecule 
of  A  comes  in  contact  with  a  molecule  of  B.  Such  collisions 
will  frequently  occur  by  reason  of  the  rapid  motion  of  both 
kinds  of  molecules.  Now  as  these  collisions  are  matters  of 
chance  it  is  very  easy  to  see  that  if  more  molecules  of  one  or 
both  kinds  are  brought  into  a  given  space  the  number  of  collisions 
of  A  molecules  with  B  molecules  will  be  increased.  On  the 
other  hand,  decreasing  the  number  of  one  or  both  kinds  of 
molecules  will  surely  decrease  the  possible  collisions  of  A  with  B 
molecules. 

Probably  not  every  collision  of  an  A  with  a  B  molecule  will 
result  in  a  union  of  the  two  to  form  A  B;  but  if,  on  the  average, 
a  certain  definite  fraction  of  the  collisions  result  in  union,  then 
we  can  say  that  the  greater  the  number  of  A  and  B  molecules 
present  in  a  given  volume,  say  i  c.c.,  of  the  gas  mixture,  the 
greater  will  be  the  number  of  A  B  molecules  formed  per  second. 
If  we  start  with  a  mixture  of  equal  numbers  of  A  and  B 


162  Introduction  to  General  Chemistry 

molecules  there  will  be  for  a  definite  pressure  and  temperature  a 
certain  number  of  AB  molecules  formed  per  second.  After 
a  short  time  the  number  of  A  and  B  molecules  will  have  de- 
creased appreciably,  so  that  now  fewer  AB  molecules  will  be 
formed  per  second,  and  as  time  goes  on,  owing  to  continual  de- 
crease in  the  numbers  of  A  and  B  molecules  present,  there  will 
be  fewer  and  fewer  AB  molecules  formed  per  second.  The 
result  will  be  that  the  rate  of  formation  of  AB  molecules  will 
be  greatest  at  the  start  and  will  gradually  decrease,  until  finally, 
if  the  reaction  is  not  reversible,  all  A  and  B  molecules  will  have 
united. 

279.  The  Kinetic  Hypothesis  Applied  to  Chemical  Equilib- 
rium.— Let  us  next  consider,  in  the  light  of  the  kinetic-molecular 
hypothesis,  the  state  of  affairs  if  a  reaction  between  gases  is 
reversible.  The  case  of  hydrogen  chloride  and  oxygen  will  serve 
as  a  good  illustration.  The  equation  is 

4HC1+O2^2C12+2H2O. 

This  reaction  takes  place  with  moderate  velocity  at  400°,  finally 
reaching  a  state  of  equilibrium  in  which  all  four  of  the  substances 
are  present. 

Suppose  we  bring  into  a  closed  vessel  at  400°  a  mixture  of 
hydrogen  chloride  and  oxygen  in  the  proportion  indicated  by  the 
equation;  that  is,  four  molecules  of  the  first  gas  to  one  of  the 
second.  The  reaction  will  begin  at  a  certain  velocity,  mole- 
cules of  hydrogen  chloride  and  oxygen  disappearing  by  uniting 
to  form  molecules  of  chlorine  and  water  vapor.  As  time  goes 
on  there  will  be  fewer  and  fewer  hydrogen  chloride  and  oxygen 
molecules  present,  so  that  the  number  of  each  uniting  per  second 
and  also  the  number  of  chlorine  and  water  molecules  formed  per 
second  will  continuously  decrease.  On  the  other  hand,  the  mole- 
cules of  chlorine  and  water  which  have  been  formed  begin  to 
reunite  to  form  hydrogen  chloride  and  oxygen.  As  the  total 
numbers  of  chlorine  and  water  molecules  present  will  increase 
as  time  goes  on,  so  the  numbers  of  these  molecules  which  react 
and  so  disappear  per  second  will  also  increase.  The  final  result 
will  be  that  in  each  second  there  will  be  just  as  many  molecules 


Chemical  Equilibrium  163 

of  chlorine  and  water  disappearing  as  the  numbers  of  each  formed. 
The  same  sort  of  thing  will  be  true  for  the  hydrogen  chloride  and 
oxygen — as  many  molecules  of  each  will  finally  be  produced  per 
second  as  the  numbers  that  disappear.  When  this  condition 
is  reached  no  further  change  in  the  number  of  any  of  the  four 
sorts  of  molecules  will  take  place,  although  chemical  change  will 
go  on  continuously.  The  system  is  then  in  a  state  of  equilibrium. 

We  may  now  take  up  a  study  of  a  number  of  reversible  reac- 
tions which  reach  a  state  of  equilibrium. 

280.  Ferric  Chloride  and  Ammonium  Sulfocyanate. — If  we 
add  to  a  very  dilute  solution  of  ferric  chloride,  FeCl3,  which  is 
faintly  yellow  in  color,  a  dilute  solution  of  ammonium  sulfo- 
cyanate,  NH4NCS,  which  is  colorless,  a  blood-red  solution  results. 
This  red  substance  is  ferric  sulfocyanate,  Fe(NCS)3,  which  is 
formed  thus: 

FeCl3+3NH4NCS->Fe(NCS)3+3NH4Cl. 

Let  us  now  consider  how  we  may  discover  whether  this 
reaction  is  complete  when  the  two  substances  on  the  left-hand 
side  of  the  equation  are  mixed  in  the  indicated  proportion  or 
whether  a  state  of  equilibrium  results.  The  experiment  may  be 
carried  out  on  the  lecture  table  in  the  following  manner: 

To  2  liters  of  water  we  add  20  c.c.  of  a  decinormal  solution  of 
ferric  chloride  and  20  c.c.  of  a  decinormal  solution  of  ammonium 
sulfocyanate,  which  is  just  the  amount  indicated  by  the  equation 
as  required  for  the  amount  of  ferric  chloride  present.  Let  us  now 
divide  the  red  solution  into  four  equal  portions,  which  we  may 
place  in  four  similar  cylinders  or  beakers.  Suppose  we  now  add 
to  the  solution  in  one  of  the  cylinders  20  c.c.  more  of  ammonium 
sulfocyanate  solution.  The  solution  will  be  seen  to  become 
deeper  red  in  color,  which  means  that  more  red  ferric  sulfo- 
cyanate has  been  formed.  Now  this  fact  may  be  explained  in 
either  of  two  ways:  first,  that  we  had  by  mistake  used,  in  the 
first  place,  less  than  the  correct  proportion  of  ammonium  sulfo- 
cyanate indicated  by  the  equation;  or,  secondly,  that  a  state  of 
equilibrium  existed  in  the  solution  and  that  the  increased  con- 
centration of  ammonium  sulfocyanate  had  shifted  the  equilibrium 
so  as  to  form  more  ferric  sulfocyanate. 


164  Introduction  to  General  Chemistry 

We  can  test  the  truth  or  falsity  of  the  first  supposition  very 
easily.  If  the  original  mixture  contained  less  than  the  correct 
proportion  of  ammonium  sulfocyanate,  then  there  would  be 
an  excess  of  ferric  chloride,  and  the  addition  of  more  of  this  salt 
would  not  increase  the  amount  of  ferric  sulfocyanate  and  so 
increase  the  depth  of  red  color.  Let  us  add,  therefore,  20  c.c. 
more  ferric  chloride  to  the  solution  in  the  second  cylinder.  It 
becomes  deeper  red!  This  seems  to  show  that  we  are  dealing 
with  a  condition  of  equilibrium  as  indicated  by  the  double  arrows 
of  the  following  equation: 

FeCl3+3NH4NCS±5Fe(NCS)3+3NH4Cl. 

If  such  is  the  case,  then  the  addition  of  ammonium  chloride  to 
the  solution  in  the  third  cylinder  should  cause  a  partial  fading 
of  the  red  color  by  reason  of  the  partial  disappearance  of  the  red 
ferric  sulfocyanate.  Now  this  is  actually  what  happens  when  the 
experiment  is  carried  out,  as  can  be  seen  by  comparison  with  the 
color  of  the  solution  in  the  fourth  cylinder. 

It  is  clear,  therefore,  that  we  have  here  a  case  of  chemical 
equilibrium  in  which  all  four  of  the  substances  represented  in  the 
equation  can  coexist  in  the  same  solution.  When  we  added  more 
ammonium  sulfocyanate  to  the  solution  in  the  first  cylinder  we 
increased  the  number  of  molecules  of  this  salt  and  so  increased 
the  chances  of  collision  of  ferric  chloride  molecules  with  am- 
monium sulfocyanate  molecules  and  this  increased  the  number 
of  ferric  sulfocyanate  molecules  formed  per  second.  This 
caused  an  increase  in  the  total  amount  of  the  latter  salt,  and 
thus  gave  rise  very  quickly  to  a  new  state  of  equilibrium  in 
which  the  proportion  of  ferric  salt  in  the  form  of  red  sulfocyanate 
was  greater  than  at  first. 

The  addition  of  more  ferric  chloride  to  the  solution  in  the 
second  cylinder  caused  a  similar  shift  of  equilibrium  for  anala- 
gous  reasons.  It  is  a  general  rule  that  increasing  the  concentration 
of  either  of  the  reacting  substances  on  the  same  side  of  an  equation 
causes  a  shift  in  equilibrium  so  as  to  form  more  of  the  substances 
on  the  other  side  of  the  equation.  This  rule  is  also  illustrated  by 
the  fact  that  when  more  ammonium  chloride  was  added  to  the 


Chemical  Equilibrium  165 

solution  in  the  third  cylinder  the  color  partially  faded;  this 
showed  that  some  of  the  red  ferric  sulfocyanate  had  disappeared, 
and  thus  indicated  that  more  ferric  chloride  and  ammonium 
sulfocyanate  had  been  formed. 

281.  Hydrogen  and  Iodine. — We  have  already  seen  (264) 
that  hydrogen  unites  with  iodine  vapor  with  appreciable  speed 
at  about  400°.  The  reaction  is  not  complete,  but  reaches  a 
state  of  equilibrium  while  there  are  still  considerable  uncom- 
bined  substances  present.  The  equation  is 


That  the  reaction  is  reversible  is  easily  shown  by  heating  hydro- 
gen iodide  gas,  when  the  purple  vapors  of  iodine  appear.  If  the 
temperature  is  370°,  equilibrium  is  reached  when  one-fifth  of  the 
hydrogen  iodide  has  dissociated  into  free  iodine  and  free  hydro- 
gen. This  means  that  out  of  every  1,000  molecules  of  hydrogen 
iodide  taken,  200  have  dissociated  and  800  remain  when  the 
state  of  equilibrium  is  reached.  The  equation  shows  that  one 
molecule  of  hydrogen  and  one  of  iodine  are  formed  by  the  disso- 
ciation of  two  molecules  of  hydrogen  iodide.  Therefore  for 
every  200  molecules  of  the  compound'  dissociated  there  would 
be  formed  100  molecules  of  hydrogen  and  100  of  iodine.  The 
equilibrium  mixture  resulting  from  every  1,000  molecules  of 
hydrogen  iodide  taken  consists,  therefore,  of  800  molecules  of 
hydrogen  iodide,  100  molecules  of  hydrogen,  and  100  molecules 
of  iodine. 

If  we  bring  together  in  a  closed  vessel  equal  numbers  of 
molecules  of  hydrogen  and  iodine  and  heat  at  370°  until  equilib- 
brium  is  reached  we  shall  find  that  for  every  500  molecules  of 
hydrogen  and  500  molecules  of  iodine  taken  there  result  800 
molecules  of  hydrogen  iodide,  100  molecules  of  hydrogen,  and 
100  of  iodine.  In  other  words,  just  the  same  proportion  as 
would  be  obtained  by  starting  with  pure  hydrogen  iodide  gas. 

282.  The  Criterion  of  Equilibrium. — In  all  cases  of  reactions 
reaching  a  condition  of  equilibrium  the  resulting  mixture  has 
the  same  proportions  of  all  substances,  whether  we  start  with 
the  substances  on  one  side  of  the  equation  or  with  equivalent 


1  66  Introduction  to  General  Chemistry 

amounts  of  those  on  the  other  side.  Therefore,  if  we  wish  to 
know  whether  a  given  reaction  has  reached  equilibrium  we  bring 
together  the  substances  which  would  be  the  products  of  the  first 
reaction.  If  the  resulting  reaction  then  gives  a  mixture  of  the 
same  composition  as  that  obtained  in  the  first  case  we  conclude 
that  both  reactions  have  reached  equilibrium. 

283.  Equilibrium  Constant.  —  In  the  hydrogen  and  iodine 
reaction 


the  rate  of  union  of  hydrogen  and  iodine,  which  we  may 
call  the  speed  of  the  reaction  from  left  to  right,  will  depend  on 
the  numbers  of  molecules  of  these  two  elements  present  in  each 
c.c.  It  would  seem  probable  that  for  a  fixed  number  of  hydrogen 
molecules  per  c.c.  the  speed  of  union  would  vary  directly  as  the 
number  of  iodine  molecules,  and  vice  versa;  so  that  this  speed 
should  be  proportional  to  the  product  of  the  number  of  hydrogen 
molecules  Nx  and  the  number  of  iodine  molecules  N2  present  in 
each  c.c.  of  the  gas  mixture.  That  is,  the  speed  of  union,  SIt  is 
proportional  to  Ni  times  N*\  or,  algebraically, 


where  £t  is  a  constant  proportionality  factor. 

On  the  other  hand,  the  reverse  change  involves  the  formation 
of  hydrogen  and  iodine  from  hydrogen  iodide,  and  we  see  by 
referring  to  the  equation  that  two  molecules  of  hydrogen  iodide 
must  react  in  order  that  one  molecule  of  hydrogen  and  one 
molecule  of  iodine  may  be  formed.  This  fact  would  make  it 
seem  necessary  for  two  molecules  of  hydrogen  iodide  to  collide 
in  order  that  the  change  could  occur.  If  so,  increasing  the  num- 
ber of  HP  molecules  in  each  c.c.  would  increase  for  each  mole- 
cule the  chances  per  second  of  collision  and,  in  fact,  doubling 
the  number  of  molecules  of  this  gas  per  c.c.  would  increase  the 
total  number  of  the  chances  per  second  fourfold,  etc.  In  other 
words,  the  number  of  collisions  per  second  of  HI  molecules  with 

1  It  has  become  customary  in  chemical  literature  to  use  formulae  of  simple 
substances  as  abbreviations  for  the  names  of  these  substances;  especially  in  cases 
of  frequent  repetition. 


Chemical  Equilibrium  167 

one  another  will  be  proportional  to  the  square  of  the  number  of 
molecules  of  this  sort  in  each  c.c.  The  details  of  the  method  of 
arriving  at  this  conclusion  need  not  be  considered  at  present. 
If  we  call  the  speed  of  change  of  hydrogen  iodide  into  hydrogen 
and  iodine  S2  and  call  the  number  of  HI  molecules  in  i  c.c.  N3, 
then  it  is  plain  that  this  speed  is  proportional  to  N32,  or 


where  k2  is  a  constant  proportionality  factor. 

Let  us  now  think  of  the  state  of  affairs  when  equilibrium  has 
resulted.  The  speed  of  formation  of  hydrogen  iodide  which 
is  equal  to  the  speed  of  union  of  hydrogen  and  iodine,  &,  is  now 
just  equal  to  the  speed  of  dissociation,  S2,  of  the  hydrogen  iodide. 
This  must  be  the  case,  as  otherwise  further  changes  in  the  pro- 
portions of  the  three  substances  would  still  be  taking  place  and 
the  mixture  would  not  be  in  equilibrium.  For  the  state  of 
equilibrium,  therefore,  we  may  write 

Sx=S2> 
and  hence 


or 

N32    =k 
N,XN2    k2' 

Now  &  and  k2  are  both  constant  quantities  for  the  reaction 
under  consideration  if  the  temperature  is  fixed,  and  therefore 
their  quotient  is  a  constant,  so  that  we  may  write 


a  constant.    Therefore, 

AT"2 
—  —  ^—  =  K 


This  algebraic  equation  means  that  for  the  condition  of  equilibrium 
at  a  fixed  temperature  the  square  of  the  number  of  molecules  per  c.c. 
of  HI  divided  by  the  product  of  the  numbers  of  molecules  of  H2  and 
1  2  is  a  fixed  or  constant  quantity.  This  matter  can  perhaps  be 
made  a  little  plainer  by  use  of  a  numerical  example.  We  have 


1  68  Introduction  to  General  Chemistry 

seen  that  at  370°  the  equilibrium  mixture  which  results  from 
i  ,000  original  HI  molecules  consists  of  800  molecules  of  HI,  100 
of  H2,  and  100  of  I2.  In  each  c.c.  of  such  an  equilibrium  mixture 
the  total  number  of  molecules  will  be  very  great  ;  but,  of  course, 
the  numbers  of  each  kind  will  be  in  the  same  proportion  as  for  a 
total  of  1,000  molecules,  and  therefore 

800* 

K=  -  —  -  =64. 
looX  100 

If  we  start  with  unequal  instead  of  equal  numbers  of  molecules 
of  hydrogen  and  iodine  we  can  calculate  by  means  of  the  equa- 
tion 


what  the  state  of  equilibrium  will  be.  For  example,  suppose  we 
start  with  a  mixture  of  hydrogen  and  iodine  containing  four 
times  as  much  hydrogen  as  would  theoretically  be  necessary 
for  the  iodine  taken;  that  is,  four  molecules  of  hydrogen  for  one 
of  iodine.  Calculation  shows  that  if  we  start  with  800  mole- 
cules of  hydrogen  and  200  molecules  of  iodine,  when  equilibrium 
is  reached,  out  of  a  total  of  1,000  molecules  392  will  be  hydrogen 
iodide,  604  will  be  free  hydrogen,  and  4  will  be  free  iodine. 

.284.  Ammonia  and  Water.  —  Several  reactions  already  studied 
reach  a  condition  of  equilibrium;  three  of  the  most  familiar  of 
these  may  now  be  considered  as  additional  examples  of  the  subject 
under  discussion.  Ammonia  gas,  NH3,  dissolves  abundantly  in 
water,  giving  a  solution  which  turns  litmus  blue  and  forms  salts 
with  acids.  The  solution  contains  ammonium  hydroxide, 
formed  by  the  union  of  ammonia  with  water  (91)  : 

NH3+H2O->NH4OH. 

The  solution  smells  strongly  of  ammonia  and  if  it  is  boiled  a 
short  time  all  of  the  gas  is  given  off.  This  shows  that  ammonium 
hydroxide  easily  dissociates  into  its  constituents.  It  seems 
highly  probable  that  in  the  water  solution  a  condition  of  equilib- 
rium exists,  as  indicated  in  the  equation 

NH3+H2O±>NH4OH, 


Chemical  Equilibrium  169 

both  free  ammonia  and  ammonium  hydroxide  being  present. 
Heating  such  a  solution  renders  the  free  ammonia  less  soluble, 
and  as  this  partially  escapes,  the  rate  of  formation  of  ammonium 
hydroxide  falls  farther  and  farther  behind  the  rate  of  dissociation 
of  this  compound  until  finally  all  of  the  latter  has  disappeared. 

285.  Carbon  Dioxide  and  Water. — A  water  solution  of  carbon 
dioxide,  C02,  contains  carbonic  acid,  H2CO3.     But  such  a  solution 
easily  gives  off  carbon  dioxide,  especially  if  warmed;  which  leads 
us  to  conclude  that  the  reaction  is  a  reversible  one,  and  that  in 
the  solution  there  is  a  state  of  equilibrium  as  represented  by 
the  equation 

C02+H20±;H2C03.  (152) 

286.  Sulfur  Dioxide  and  Water. — Sulfur  dioxide,  S02,  which 
is  formed  when  sulfur  burns,  is  a  colorless  gas  with  a  suffocating 

odor: 

S+O2->SO2. 

It  is  easily  soluble  in  water,  giving  a  solution  which  smells 
strongly  of  the  gas  and  has  acid  properties.  The  solution  con- 
tains a  compound,  sulfurous  acid,  H2S03.  This  solution  gives  off 
all  of  its  sulfur  dioxide  when  boiled,  and  we  conclude,  therefore, 
that  the  acid  easily  decomposes  into  its  constituents,  water  and 
sulfur  dioxide,  and  that  in  the  solution  we  have  a  state  of  equilib- 
brium,  as  represented  by  the  equation 

SO2+H2O^H2SO3. 

287.  The  Effect  of  Pressure  on  a  System  in  Equilibrium. — 

Suppose  we  have,  say,  i  liter  of  water  saturated 
with  a  gas,  say  oxygen,  at  a  fixed  temperature 
and  at  one-atmosphere  pressure.  To  say  that 
the  water  is  saturated  with  the  gas  means  that 
a  condition  of  equilibrium  exists  between  solu- 
tion and  gas.  Let  us  suppose  the  solution  and 
gas  are  contained  in  a  cylinder  fitted  with  a  gas- 
tight  piston  (Fig.  38)  and  that  the  volume  of  the 


undissolved  oxygen  gas  above  the  solution  is  i  FIG.  38 

liter.    If  now  we  double  the  pressure  on  the  gas 

more  of  the  gas  passes  into  solution,  finally  producing  a  new 


170  Introduction  to  General  Chemistry 

state  of  equilibrium.  By  reason  of  the  fact  that  part  of  the 
gas  dissolved  when  the  pressure  was  doubled  the  volume  of  the 
remaining  gas  will  not  be  half  a  liter,  as  we  should  expect  if  no 
additional  quantity  of  oxygen  dissolved  in  the  water  present, 
but  appreciably  less  than  half  a  liter.  The  effect,  therefore,  of 
increasing  the  pressure  on  the  system  is  to  cause  its  volume  to 
become  smaller  than  would  be  the  case  if  no  shift  of  equilibrium 
had  occurred.  This  is  the  way  in  which  an  increase  of  pressure 
always  affects  a  system  in  equilibrium:  the  state  of  equilibrium 
shifts  in  such  a  way  as  to  cause  a  greater  decrease  in  volume  than 
would  be  the  case  if  no  change  in  the  state  of  equilibrium  occurred. 
Let  us  consider  another  case.  We  may  inquire  how  the 
equilibrium  represented  by  the  equation 


would  be  affected  by  increase  of  pressure.  We  see  by  reference 
to  the  equation  that  four  volumes  of  HC1  and  one  of  O2  give  two 
volumes  of  C12  and  two  of  H20;  that  is,  that  when  the  reaction 
takes  place  from  left  to  right  there  is  a  decrease  in  volume  from 
5  to  4.  We  should  expect,  therefore,  that  by  increasing  the 
pressure  the  equilibrium  would  shift  somewhat  from  left  to 
right;  that  is,  that  more  chlorine  and  water  would  be  formed 
at  the  expense  of  the  hydrogen  chloride  and  oxygen;  and  this 
is  exactly  what  actually  happens. 

The  effect  of  increase  of  pressure  on  any  system  in  equilibrium 
is,  in  all  cases,  to  shift  the  equilibrium  so  as  to  favor  the  formation 
of  substances  occupying  a  smaller  wlume.  In  case  no  change  of 
volume  accompanies  a  chemical  reaction,  then  the  state  of 
equilibrium  is  not  affected  by  change  of  pressure.  The  reaction 


is  an  example  of  this  sort.  Here  one  volume  of  hydrogen  and 
one  volume  of  iodine  vapor  react  to  form  two  volumes  of  hydro- 
gen iodide,  so  that  no  change  of  volume  occurs  when  the  reaction 
takes  place.  It  has  been  found  by  careful  investigation  that 
the  equilibrium  proportion  of  the  three  substances  is  not  changed 


Chemical  Equilibrium  171 

by  altering  the  pressure,  as  long  as  the  temperature  remains 
constant. 

288.  Effect  of  Temperature  on  a  System  in  Equilibrium. — We 
have  already  learned  (112)  that  the  vapor  pressure  of  water 
increases  with  increase  of  temperature.  We  know  also  that  a 
large  amount  of  heat  is  absorbed  when  water  is  evaporated;  at 
1 00°  it  requires  540  calories  to  change  one  gram  of  water  into 
steam.  This  is  the  so-called  latent  heat  of  vaporization.  If 
we  have,  in  a  closed  vessel,  water  in  equilibrium  with  its  vapor, 
and  then  apply  heat,  two  effects  are  produced:  the  temperature 
is  raised  and  the  vapor  pressure  is  increased.  The  increase  in 
pressure  is  caused  by  the  evaporation  of  some  water,  and  this 
evaporation  absorbs  some  of  the  heat  which  has  been  applied. 
This  is  a  typical  case,  for  we  always  find  that  when  we  apply  heat 
to  any  system  in  equilibrium  that  the  state  of  equilibrium  shifts 
in  such  a  way  that  heat  is  absorbed  in  the  change.  As  heat  is 
absorbed  when  water  evaporates,  heating  causes  increased  vapor 
pressure. 

The  effect  of  temperature  on  the  solubility  of  substances 
has  already  been  studied  (134).  We  have  learned  that  heat  is 
either  absorbed  or  produced  when  a  substance  dissolves;  this  is 
the  so-called  heat  of  solution.  Substances  which  dissolve  with 
absorption  of  heat  become  more  soluble  with  rise  of  tempera- 
ture, while  those  which  dissolve  with  evolution  of  heat,  like 
anhydrous  sodium  sulfate,  Na2S04,  decrease  in  solubility  as  the 
temperature  is  raised  (134,  Fig.  27).  If  a  substance  like  the 
last  named  dissolves  with  evolution  of  heat,  its  crystallization 
out  of  a  solution  is  accompanied  by  absorption  of  heat.  In 
every  case  raising  the  temperature  causes  that  change  of  solubility 
to  occur  which  invokes  an  absorption  of  heat. 

The  state  of  chemical  equilibrium  is  shifted  in  all  cases  by  a 
change  of  temperature.  Now  we  find  that  every  chemical  reac- 
tion either  gives  out  or  absorbs  heat.  When  substances  burn,  the 
heat  given  out  is  very  great.  In  many  other  reactions  the  heat 
produced  is  considerable,  while  in  still  others  an  absorption  of 
heat  occurs.  If  a  reaction  is  reversible  (all  reactions  that  reach 
a  state  of  equilibrium  are,  of  course,  of  this  class)  and  produces 


172  Introduction  to  General  Chemistry 

heat  when  it  goes  in  one  direction,  it  absorbs  an  equal  amount  of 
heat  for  the  same  quantity  of  materials  transformed  when  it  goes 
in  the  opposite  direction. 
If  the  reaction 

(264) 


has  reached  a  state  of  equilibrium  at  370°,  out  of  every  1,000 
molecules  present  800  will  be  HI,  100  H2,  and  100  I2.  If  the 
temperature  is  then  raised  to  440°  and  held  constant  until  a 
new  state  of  equilibrium  is  reached,  the  gas  mixture  will  consist 
of  780  molecules  of  HI,  no  of  H2,  and  no  of  I2.  Part  of  the 
HI  has  changed  to  H2  and  I2,  and  the  equilibrium  may  be  said 
to  have  shifted  from  right  to  left.  At  temperatures  between 
370°  and  440°  the  change  of  HI  into  H2  and  I2  takes  place  with 
absorption  of  heat.  We  see,  then,  that  raising  the  temperature 
causes  the  equilibrium  to  shift  in  the  direction  that  involves  an 
absorption  of  heat.  Now  this  is  a  perfectly  general  law  for 
chemical  changes,  just  as  it  is  also  for  physical  changes  like  the 
vaporizing  of  a  liquid  and  dissolving  of  a  solid. 

289.  The  Effect  of  Removing  One  Product  of  a  Reaction.  — 
The  reaction  represented  by  the  equation 

NaCl+H2SO4±5NaHSO4+HCl  (103,  253) 

has  already  been  studied  rather  fully.  We  may  summarize 
the  facts  briefly,  as  follows:  The  action  of  concentrated  sulfuric 
acid  on  dry  salt  gives  sodium  hydrogen  sulfate,  NaHSO4,  and 
hydrogen  chloride  gas,  the  reaction  going  nearly  to  completion 
in  the  direction  of  the  lower  arrow  in  the  equation  given  above 
if  the  mixture  is  warmed.  On  the  other  hand,  if  a  cold  saturated 
solution  of  sodium  hydrogen  sulfate  is  mixed  with  concentrated 
hydrochloric  acid  —  that  is,  a  saturated  solution  of  hydrogen 
chloride  in  water  —  an  abundant  precipitate  of  solid  salt,  NaCl,  is 
formed.  This  reaction  is,  we  see,  just  the  reverse  of  the  other. 
If  now  we  mix  a  dilute  solution  of  salt  with  dilute  sulfuric  acid, 
we  see  no  visible  change.  We  also  see  no  change  upon  mixing 
a  dilute  solution  of  sodium  hydrogen  sulfate  with  dilute  hydro- 
chloric acid. 


Chemical  Equilibrium  173 

We  are  now  in  position  to  explain  all  the  facts  of  the  foregoing 
paragraph  from  the  standpoint  of  chemical  equilibrium.  If 
we  bring  together  dilute  solutions  of  either  pair  of  substances 
in  the  reaction 

NaCl+H2$O4^NaHSO4+HCl, 

the  resulting  solution  probably  contains  all  four  substances,  side 
by  side,  in  a  state  of  equilibrium.  But  we  cannot  notice  any 
effect  of  the  mixing,  because  in  the  presence  of  much  water  all 
four  are  held  completely  in  solution,  since  all  four  are  more  or 
less  readily  soluble  in  water.  If,  however,  but  little  water  is 
present,  the  least  soluble  of  the  four  substances,  common  salt, 
may  partially  separate.  This  is  the  case  when  a  concentrated 
solution  of  NaHS04  is  mixed  with  concentrated  HC1.  The 
reason  is  a  simple  one:  the  substances  taken  react  partially  to 
form  NaCl  and  H2S04  in  the  sense  of  the  upper  arrow;  but  the 
amount  of  NaCl  so  formed  is  more  than  the  water  present  can 
hold  in  solution;  so  the  excess  NaCl  separates  out  in  the  solid 
form.  This  separation  of  NaCl  continues  until  the  four  sub- 
stances in  the  solution  have  reached  amounts  which  can  and  do 
exist  in  equilibrium  with  one  another.  Removing  the  solid 
NaCl  which  has  separated,  or  adding  more  solid  salt,  will  in  no 
way  alter  the  amounts  of  any  of  the  four  substances  contained  in 
the  solution. 

When  concentrated  H2SO4  is  mixed  with  dry  NaCl,  NaHS04 
and  HC1  begin  to  be  formed.  Now  HC1  is  but  slightly  soluble 
in  concentrated  H2S04  and,  being  a  gas,  it  at  once  escapes  from 
the  mixture.  Warming  the  mixture  also  promotes  the  escape 
of  the  HC1,  since  the  higher  the  temperature  the  smaller  the 
solubility  of  the  gas  in  the  concentrated  H2S04.  The  escape  of 
the  HC1  gas  also  has  another  fundamental  effect  on  the  reaction. 
In  order  that  any  reaction  may  reach  a  state  of  equilibrium  it 
must  be  reversible;  but  this  reaction  cannot  go  in  the  reverse 
direction  if  the  HC1  escapes  from  the  reacting  mixture  as  fast 
as  it  is  formed.  The  result  is  that  if  no  water  is  present,  con- 
centrated H2SO4  and  dry  NaCl  react  practically  completely, 
giving  solid  NaHSO4  and  HC1  gas. 


174  Introduction  to  General  Chemistry 

290.  The  Action  of  Steam  on  Iron  and  the  Reverse  Action.  — 
When  steam  is  passed  over  heated  iron  (29,  Fig.  16)  hydrogen 
and  an  oxide  of  iron  are  formed.  On  the  other  hand,  if  hydrogen 
is  passed  over  the  heated  oxide,  Fe3O4,  the  products  are  iron  and 
water.  The  equation  for  these  two  reactions,  of  which  one  is 
the  reverse  of  the  other,  is 


If  we  bring  together  either  pair  of  substances  in  a  closed 
vessel  and  heat  them  for  some  time,  a  state  of  equilibrium  is 
reached  in  which  all  four  substances  are  present.  The  iron  and 
iron  oxide  are  solids,  while  the  water,  as  steam,  is  a  gas.  For 
the  condition  of  equilibrium  the  relative  amounts  of  steam  and 
hydrogen  are  always  the  same  for  a  given  temperature,  no  matter 
what  proportions  of  either  pair  of  substances  have  been  used. 
This  is  the  state  of  affairs  if  the  reaction  occurs  in  a  closed  vessel. 
But  the  results  are  entirely  different  if  the  reactions  take  place 
in  a  tube  holding  the  solids,  through  which  either  steam  in  the 
one  case  or  hydrogen  in  the  other  is  passed.  If  steam  is  passed 
through  a  tube  containing  iron,  then  the  hydrogen  which  is 
formed  is  carried  along  with  the  excess  of  steam  and  has  no 
chance  to  act  on  the  iron  oxide  which  has  been  formed.  There 
is  therefore  no  chance  for  iron  to  be  formed  again,  once  it  has 
been  changed  to  iron  oxide.  As  long  as  unchanged  iron  remains 
and  the  current  of  steam  is  continued,  the  reaction  from  left  to 
right  continues.  The  inevitable  result  is  the  complete  change 
of  the  iron  to  the  oxide.  On  the  other  hand,  if  a  current  of 
hydrogen  is  passed  over  heated  iron  oxide  contained  in  a  tube, 
the  substances  react  in  the  direction  from  right  to  left  of  the 
equation.  The  steam  which  is  formed  passes  along  with  the 
excess  of  hydrogen,  and  once  having  left  the  tube  cannot  pos- 
sibly act  on  the  iron  to  convert  it  back  into  oxide,  so  that  this 
change  also  continues  as  long  as  the  stream  of  hydrogen  is  kept 
up  and  comes  to  an  end  only  when  all  of  the  iron  oxide  has  been 
reduced  to  metallic  iron. 

291.  Conclusions.  —  We  see,  therefore,  that  a  chemical 
reaction  like  the  one  just  discussed  may  reach  a  state  of  equilib- 


Chemical  Equilibrium  175 

rium,  if  the  reverse  reaction  tends  to  take  place  noticeably,  and 
if  none  of  the  substances  involved  escape  from  the  vessel  in 
which  the  change  takes  place;  or  it  may  go  to  completion  in  one 
direction  or  the  other  if  one  of  the  products  of  either  reaction  is 
allowed  to  escape  from  the  scene  of  action. 

Whether  a  given  reaction  reaches  a  state  of  equilibrium  or 
goes  to  completion  in  one  direction  or  the  other  often  depends 
upon  the  conditions.  In  the  preparation  or  manufacture  of 
chemical  substances  it  is  usually  very  important  to  cause  equi- 
librium reactions  to  take  place  as  completely  as  possible  in 
order  to  obtain  the  maximum  possible  yields  of  the  desired 
products. 


CHAPTER  XIV 
HYDROGEN  AND  OXYGEN 

292.  Hydrogen. — Hydrogen  was  first  recognized  in  1766  as 
a  distinct  substance  by  Cavendish,  a  celebrated  English  chemist, 
who  called  the  gas  inflammable  air  and  prepared  it  by  the  action 
of  acids  on  metals.     It  was  not  until  ten  years  after  Cavendish's 
discovery  that  Lavoisier  explained  the  role  played  by  oxygen 
in  combustion  and  stated  the  law  of  the  indestructibility  of 
matter  (21)  and  thus  laid  the  foundation  for  the  doctrine  of  the 
elements  in  its  present  form.     For  this  reason  the  classification 
of  hydrogen  as  an  element  was  not  possible  at  the  time  of  its  dis- 
covery.    In  1781  Cavendish  showed  that  nothing  but  water  is 
formed  when  hydrogen  burns  and  thus  proved  that  water  is  a 
compound  of  hydrogen  and  oxygen.     The  name  hydrogen  means 
water-former. 

The  element  occurs  in  but  minute  amounts  in  the  free  form 
in  nature.  Water  is  its  most  abundant  compound;  but  it  is  also 
a  constituent  of  all  dry  animal  and  vegetable  tissues,  forming 
therein  principally  compounds  with  carbon,  oxygen,  and  nitrogen. 
Petroleum  and  natural  gas  are  compounds  of  hydrogen  with 
carbon;  coal  also  contains  considerable  combined  hydrogen. 

293.  Preparation  of  Hydrogen. — We  have  already  learned 
several  methods  by  which  free  hydrogen  can  be  obtained.     These 
may  now  be  briefly  reviewed.     Hydrogen  is  formed : 

1.  By  the  electrolysis  of  water  (27), 

2H2O->2H2+O2;  * 

2.  By  the  action  of  water  on  some  metals,  as  by  (a)  the  burn- 
ing of  magnesium  wire  in  steam  (28,  Fig.  15), 

Mg+H20>MgO+H2, 

(b)  the  passage  of  steam  over  heated  iron  turnings  (29,  Fig.  16) , 

3Fe+4H20->Fe304+4H2, 

176 


Hydrogen  and  Oxygen 


177 


(c)  the  action  of  sodium  or  potassium  on  water  (40,  86,  Table  VI, 

106), 

2Na+2H2O-»2NaOH+H2 
2K+2H2O->2KOH+H2; 

3.  By  the  action  of  hydrochloric  or  sulfuric  acid  on  zinc, 
magnesium,  iron,  or  aluminum,  as  well  as  on  several  other  metals, 


Zn+2HCl->ZnCl2+H2 
Fe+H2SO4->FeSO4+H2. 


(i49) 
d73) 


294.  Making  Hydrogen  in  the  Laboratory. — The  best  labora- 
tory method  of  making  hydrogen  consists  in  treating  zinc  with 
hydrochloric  acid  in  some  form  of 
specially  constructed  gas  generator. 
The  Kipp  apparatus,  Fig.  39,  is  the 
form  most  extensively  used.  The 
solution  used  is  made  from  equal 
volumes  of  concentrated  hydrochloric 
acid  and  water.  The  action  of  this 
generator  is  very  simple  in  principle. 
Upon  opening  the  stopcock  gas 
escapes  and  allows  the  acid  to  rise 
into  the  middle  compartment,  where 
it  acts  upon  the  zinc  and  so  produces 
a  steady  flow  of  hydrogen.  When 
the  cock  is  closed  the  gas  formed 
forces  the  acid  downward  and  causes  it  to  flow  from  the 
lower  into  the  upper  compartment.  As  soon  as  the  acid  is  out 
of  contact  with  the  zinc  all  action  stops,  and  no  more  gas  is 
produced  until  the  cock  is  again  opened.  The  Kipp  apparatus 
has  one  unfortunate  defect:  since  there  is  but  little  circulation 
of  the  solution  the  acid  in  contact  with  the  zinc  is  soon 
exhausted,  causing  the  action  to  stop  while  there  is  still  a 
large  supply  of  almost  unchanged  acid  in  other  parts  of  the 
apparatus.  To  start  the  action  again  it  is  necessary  to  empty 
all  the  solution  and  refill  with  fresh  acid;  much  acid  is  thus 
wastedi 


FIG.  39 


Introduction  to  General  Chemistry 


The  McCoy  apparatus,  shown  in  Fig.  40,  has  several  ad- 
vantages over  the  Kipp  apparatus.  The  lowest  compartment 
is  filled  as  full  as  possible  with  granulated  or  stick  zinc,  on  which 
hydrochloric  acid  drops  at  just  the  rate  required  to  keep  up  the 
stream  of  hydrogen  that  is  being  drawn  from  the  apparatus. 
When  the  stopcock  is  closed  the  gas  which  is  formed  from  the 
small  excess  of  acid  in  the  zinc  compartment  forces  the  acid  from 
the  middle  to  the  upper  compartment  and  thus  stops  the  further 
flow  of  acid  upon  the  zinc.  This  appa- 
ratus is  also  conveniently  used  for 
generating  other  gases. 

295.  The  Electrolysis  of  Water.— 
Compared  with  metals,  pure  water  is  a 
very  poor  conductor  of  electricity.  The 
addition  of  a  little  sulfuric  acid  increases 
the  electrical  conductivity  of  water 
enormously.  The  sulfuric  acid  so  added 
is  not  permanently  destroyed  in  the 
course  of  the  electrolysis  (27),  so  that 
very  little  will  serve  to  promote  the 
electrolysis  of  a  large  amount  of  water. 
The  exact  way  in  which  the  acid  behaves  will  be  discussed 
later.  Ordinarily,  poles  or  electrodes  of  the  elementary  metal 
platinum  are  employed,  since  most  other  metals  would  be 
attacked  chemically.  The  electrode  at  which  the  hydrogen  is 
liberated  is  called  the  negative  electrode,  or  cathode ;  the  other, 
at  which  the  oxygen  appears,  is  the  positive  electrode,  or 
anode. 

One  of  the  important  technical  methods  of  making  hydrogen, 
which  yields  at  the  same  time  oxygen,  consists  in  the  electrolysis 
of  water  in  which  sodium  hydroxide  is  dissolved  to  make  it  a 
good  conductor.  Here  the  cathode  is  of  iron  and  the  anode  of 
carbon.  Hydrogen  is  also  obtained  in  commercial  quantities 
as  a  by-product  in  the  manufacture  of  caustic  soda  by  the  elec- 
trolysis of  a  solution  of  common  salt. 

Hydrogen  is  often  made  for  use  in  balloons  by  the  action  of 
dilute  sulfuric  acid  on  scrap  iron. 


FIG.  40 


Hydrogen  and  Oxygen  179 

296.  The  Physical  Properties  of  Hydrogen. — We  have  already 
learned  that  hydrogen  is  colorless ;  when  perfectly  pure  it  is  also 
odorless  and  tasteless.     One  liter  of  the  gas  at  o°  and  76  cm.  pres- 
sure weighs  0.0899  g.;   and  22.4  liters,  2  g.  approximately.     It 
is  the  lightest  of  all  gases.     It  can  be  liquefied,  giving  a  colorless 
liquid  which  boils  at  —  253°,  or  only  20°  above  absolute  zero.    At  a 
somewhat  lower  temperature  the  liquid  freezes  to  a  colorless 
solid. 

Hydrogen  is  but  slightly  soluble  in  water:  100  c.c.  of  water 
dissolves  about  2  c.c.  of  the  gas  at  room  temperature. 

The  speed  of  diffusion  of  hydrogen  is  greater  than  that  of 
any  other  gas  (191). 

297.  The   Chemical   Properties   of   Hydrogen. — The   most 
important  chemical  properties  of  hydrogen  have  already  been 
studied,  but  may  now  be  briefly  reviewed.    Hydrogen  burns  with 
an  almost  non-luminous  flame,  which  is,  however,  very  much 
hotter  than  that  obtained  from  ordinary  fuel  or  illuminating  gas. 
Water  is  the  product  of  the  reaction.    Hydrogen  reacts  readily 
with  hot  copper  oxide,  forming  water  arid  copper: 

H3+ CuO^H3O+  Cu.  (33) 

Hydrogen  also  acts  on  other  metallic  oxides  at  a  red  heat,  for 

example : 

4H2+Fe3O4->4H2O+3Fe.  (290) 

Hydrogen  and  chlorine,  if  mixed  in  equal  volumes  and  ignited, 
or  exposed  to  a  bright  light,  unite  with  explosive  violence,  forming 

hydrogen  chloride, 

H2+C12->2HC1.  (243) 

A  jet  of  burning  hydrogen  lowered  into  a  jar  of  chlorine  con- 
tinues to  burn  by  reason  of  the  union  of  the  two  elements  (244). 
Hydrogen  unites  with  bromine  to  form  hydrogen  bromide 
(256)  and  with  iodine  to  form  hydrogen  iodide  (264). 

298.  The  Union  of  Hydrogen  and  Nitrogen. — A  mixture  of 
hydrogen  and  nitrogen  does  not  react  at  all  under  ordinary 
conditions.     If  electric  sparks  are  passed  through  the  mixed 


180  Introduction  to  General  Chemistry 

gases  contained  in  a  eudiometer,  Fig.  41,  a  small  amount  of 
ammonia  is  formed: 

3H2+N2-»2NH3. 

The  reaction  soon  reaches  a  state  of  equilibrium,  because  under 
the  same  conditions  ammonia  is  very  largely  decomposed  into 
its  elements.  As  the  result  of  the  reverse  reaction,  a  state  of 
equilibrium  is  reached  when  less  than  i  per  cent  of  the  ele- 
mentary gases  has  been  converted  into  ammonia.  If  the 
ammonia  is  absorbed  in  some  suitable  way, 
as  by  union  with  sulfuric  acid,  as  fast  as  it  is 
formed,  then,  with  continued  sparking,  the 
formation  of  ammonia  goes  on  until  all 
the  hydrogen  and  nitrogen  have  united.  The 
practical  method  of  making  ammonia  by  this 
reaction  will  be  discussed  in  chapter  xxi. 

299.  Heat    of    Reaction    and   Flame 
Temperature. — When  i  g.  of  hydrogen  burns, 
about  34,000  calories  of  heat  are  produced; 
FIG  ^  this  is  sufficient  to  heat  340  c.c.  of  water 

from  o°  to  100°.  By  reason  of  the  great 
amount  of  heat  produced,  the  flame  of  hydrogen  burning  in  air 
has  a  very  high  temperature.  When  hydrogen  burns  in  pure 
oxygen  instead  of  in  air,  the  flame  is  much  hotter,  but  not 
because  a  greater  amount  of  heat  is  produced  by  the  burning 
of  a  given  amount  of  hydrogen,  since  the  quantity  of  heat  is 
the  same  in  the  two  cases.  When  hydrogen  burns  in  air,  the 
nitrogen,  which  forms  four-fifths  by  volume  of  the  air,  is 
heated  to  the  flame  temperature  along  with  the  steam  formed. 
But  in  pure  oxygen  no  nitrogen  is  present,  and  so  the  tempera- 
ture reached  by  the  flame  is  much  higher,  as  there  is  far  less 
material  to  be  heated. 

300.  The  Oxyhydrogen  Blowpipe. — Fig.  42  shows  an  oxy hy- 
drogen blowpipe.  The  two  gases  mix  in  the  proper  proportions 
before  issuing  from  the  jet.  The  temperature  of  the  flame  is  high 
enough  to  melt  platinum,  which  cannot  be  melted  in  a  Bunsen 
flame  supplied  with  fuel  or  illuminating  gas. 


Hydrogen  and  Oxygen  181 

301.  The  Limelight. — A  very  bright  light  is  produced  when 
an  oxyhydrogen  flame  strikes  a  stick  of  quicklime,  by  reason  of 
the  bright  white  heat  to  which  the  lime  is  raised.    This  is  the 
so-called  limelight,  which  was  very  extensively  used  before  the 
electric  arc  light  was  perfected  and  which  is   still  frequently 
used  in  rural  communities.     In  place  of  lime  other  difficultly 
fusible   white   oxides   may   be    employed.     For   this   purpose 
thorium  oxide  containing  i  per  cent  of  cerium  oxide  is  much 
superior  to  lime. 

302.  Ignition  Temperature. — A  mixture  of  hydrogen  and 
oxygen  in  their  combining  proportion  remains  unchanged  for  any 


FIG.  42 

length  of  time  at  room  temperature,  but  if  brought  in  contact 
with  a  flame  or  electric  spark  it  explodes  violently.  The  explo- 
sion is  due  to  the  great  increase  in  volume  of  the  reaction  product, 
steam,  caused  by  the  almost  instantaneous  reaction,  with  its 
attendant  heat  production.  But  we  may  inquire  why  a  reaction 
which  does  not  take  place  at  room  temperature  can  become  explo- 
sive. Investigation  shows  that  a  mixture  of  the  two  gases  reacts 
perceptibly  at  450°,  and  that  the  formation  of  water  goes  on 
faster  the  higher  the  temperature,  but  that  the  mixture  does  not 
become  explosive  until  the  temperature  reaches  about  600°.  It 
is  easy  to  see  why  explosion  finally  occurs  when  the  temperature 
is  raised.  While  the  reaction  is  taking  place  slowly,  heat  is  being 
produced  by  the  chemical  change;  below  600°  the  rate  of  change 
is  so  slow  that  heat  is  lost  by  the  gas  mixture  faster  than  it  is 
produced.  Above  600°  the  reaction  goes  faster,  so  that  heat  is 
produced  more  rapidly  than  it  is  lost,  and  this  causes  the  gas 
mixture  to  grow  hotter;  and  the  hotter  it  gets  the  faster  the 
reaction  goes,  until  soon  it  proceeds  with  enormous  rapidity,  and 
this  constitutes  an  explosion.  For  any  combustible  substance 
there  is  some  temperature  to  which  it  must  be  heated  before  its 


1 82  Introduction  to  General  Chemistry 

rate  of  production  of  heat  by  union  with  oxygen  exceeds  its  rate 
of  loss  of  heat;  if  heated  to  this  temperature  the  substance  takes 
fire  and  continues  to  burn.  This  point  is  called  the  ignition 
temperature. 

Hydrogen,  issuing  from  a  jet,  burns  quietly  in  air  when 
ignited.  This  is  because  the  actual  union  with  oxygen  can  occur 
only  as  fast  as  the  two  gases  can  reach  one  another  by  diffusion 
(191),  one  from  the  jet,  the  other  from  the  surrounding  air. 
The  flame  is  the  reacting  gas  mixture,  which  is  raised  to  incan- 
descence by  the  great  heat  produced  by  the  union. 

303.  Platinum  as  a  Catalytic  Agent. — The  element  platinum 
can  be  deposited  on  asbestos  as  a  thin,  spongy  coating  by  dipping 
a  bit  of  fibrous  asbestos  in  a  solution  of  platinum  chloride, 
PtCl4,  drying  the  material  and  holding  it  in  a  Bunsen  flame  for 
a  minute.    The  salt  decomposes  into  its  elements  thus : 

PtCl4->Pt+2Cl2. 

If  a  jet  of  cold  hydrogen  gas  is  directed  on  the  cold  platinized 
asbestos,  the  latter  gets  red-hot  and  sets  fire  to  the  hydrogen. 
Spongy  platinum  absorbs  gases  to  a  marked  extent.  Heat  is 
produced  in  this  way,  and  this  ignites  the  intimate  mixture  of 
hydrogen  and  oxygen  condensed  on  the  surface  of  the  metal. 
The  platinum  itself  is  entirely  unchanged  and  will  continue 
active  in  this  way  indefinitely.  A  substance  which  initiates  or 
promotes  a  chemical  reaction  without  itself  being  changed  is 
called  a  catalytic  agent  (239). 

304.  The  Use  of  Hydrogen  in  Balloons. — The  lifting  power 
of  a  balloon  filled  with  hydrogen  can  easily  be  calculated.    Since 
i  liter  of  air  under  standard  conditions  weighs  i .  29  g.  and 
i  liter  of  hydrogen  weighs  0.09  g.,  the  difference,  i .  2  g.,  repre- 
sents the  lifting  power  per  liter  of  capacity  of  a  balloon.    At 
higher  temperature  and  lower  pressure  the  lifting  power  is 
smaller.    If  a  Zeppelin  has  a  capacity  of  5,000,000  liters  its 
lifting  power  will  be  about  6,000  kilos,  or  more  than  13,000 
pounds. 

305.  Oxygen. — Oxygen  in  the  form  of  compounds  makes  up 
about  one-half  by  weight  of  the  matter  forming  the  crust  of  the 


Hydrogen  and  Oxygen  183 

earth.  It  also  constitutes  89  per  cent  by  weight  of  water  and 
21  per  cent  by  volume  of  the  air.  The  oxygen  of  the  air  is  not 
chemically  combined  but  is  only  mixed  with  nitrogen  and  small 
amounts  of  other  gases  present.  Oxygen  was  first  prepared  by 
Priestley,  in  England  (1774),  by  heating  mercuric  oxide, 

2HgO-»2Hg+02. 

At  practically  the  same  time  Scheele,  in  Sweden,  made  oxygen 
by  this  method  and  also  by  heating  potassium  nitrate, 

2KN03->2KNOa-fOa. 

The  salt  KNOa  is  called  potassium  nitrite.  The  name  oxygen, 
which  means  acid-former,  was  given  to  the  gas  by  Lavoisier,  who 
believed  that  it  was  a  necessary  constituent  of  all  acids.  At 
that  time  hydrochloric  acid,  which  was  called  muriatic  or  marine 
acid,  was  thought  to  contain  oxygen.  We  now  know  many 
other  acids  which  do  not  contain  oxygen. 

306.  The  Preparation  of  Oxygen. — We  have  already  learned 
several  ways  by  which  oxygen  may  be  made.  The  heating  of 
mercuric  oxide  and  the  electrolysis  of  water  (14, 295)  have  already 
been  fully  studied.  We  have  also  seen  (245)  that  oxygen  is 
formed  when  chlorine  water  is  exposed  to  sunlight: 

2C13+2H2O->O2-HHC1. 

The  heating  of  certain  salts  which  are  rich  in  oxygen  is  also 
a  simple  way  of  making  the  gas.  The  behavior  of  potassium 
nitrate,  KN03,  is  given  in  the  preceding  paragraph.  Potassium 
chlorate,  KC1O3,  is  easily  decomposed  by  heat  according  to  the 
following  equation: 

2KC103-*2KC1+302. 

This  last  reaction  is  the  one  usually  employed  in  making  small 
amounts  of  oxygen  in  the  laboratory.  It  can  be  carried  out  in 
a  test  tube,  a  small  flask,  or  a  retort.  The  crystals  of  potassium 
chlorate  first  melt,  and  at  a  little  higher  temperature  the  liquid 
seems  to  boil,  by  reason  of  the  oxygen  given  off. 


184  Introduction  to  General  Chemistry 

The  change  of  the  chlorate,  KC103,  into  chloride,  KC1,  does 
not  take  place  completely  in  one  step.  The  first  stage  of  the 
reaction  is  probably  represented  by  the  equation, 

ioKC103->4KCl+6KClO4+3O2. 

The  salt  KC104,  called  potassium  perchlorate,  can  also  be  decom- 
posed by  heat,  thus: 

KC104->KC1+2O2. 

This  last  reaction  requires  a  higher  temperature  than  the  first. 
If  the  heating  of  the  chlorate  is  stopped  when  about  one-fifth  of 
its  total  oxygen  has  been  given  off,  KC1O4  will  be  found  in  the 
residue. 

In  making  oxygen  from  potassium  chlorate  two  precautions 
should  be  observed :  first,  the  material  must  be  entirely  free  from 
bits  of  wood,  paper,  etc.,  which  are  easily  combustible;  and,  sec- 
ondly, the  heating  must  be  gentle,  as  otherwise  the  decomposition 
may  occur  explosively. 

When  powdered  potassium  chlorate  is  mixed  with  about  half 
its  weight  of  manganese  dioxide,  Mn02,  it  will  give  off  its 
oxygen  rapidly  at  a  temperature  far  below  that  at  which  the 
pure  chlorate  starts  to  decompose.  Since  manganese  dioxide 
alone  does  not  give  off  any  of  its  oxygen  until  a  rather  high 
temperature  is  reached,  and  is  not  changed  itself  in  promoting 
the  decomposition  of  the  potassium  chlorate,  we  must  consider 
that  the  former  substance  acts  only  as  a  catalytic  agent  in  pro- 
moting the  decomposition  of  potassium  chlorate. 

307.  Oxygen  from  Sodium  Peroxide. — Sodium  peroxide, 
Na202,  is  a  solid  made  by  burning  metallic  sodium, 

2Na+O2-»Na2O2. 

The  trade  name  of  the  material  is  oxone;  it  is  supplied  in  the 
form  of  lumps  or  sticks.  Water  acts  on  it  as  follows: 

2Na2O2+  2H2O-»4NaOH+O2. 

By  dropping  water  on  lumps  of  oxone  contained  in  a  suitable 
apparatus  (236,  Fig.  30)  a  steady  stream  of  oxygen  is  obtained. 


Hydrogen  and  Oxygen  185 

The  method  is  rather  expensive,  but  it  is  very  convenient,  since 
the  action  stops  when  the  supply  of  water  is  turned  off  and  can 
be  started  again  at  will. 

308.  Oxygen  from  Other  Oxides.  —  Lead  dioxide,  Pb02,  when 
strongly  heated  gives  oxygen  and  lead  monoxide,  or  litharge, 

PbO: 

2PbO2->2PbO+O2. 

Manganese  dioxide  is  also  decomposed  at  a  high  temperature, 
thus: 

3MnO2->Mn3O4+O2. 

309.  Technical   Methods   of   Making    Oxygen.  —  The   elec- 
trolysis of  water  is  an  important  technical  method  of  making 
oxygen.     It  also  yields  hydrogen  and  has  already  been  described. 
By  far  the  larger  part  of  the  oxygen  of  commerce  is  made  from 
liquid  air.     This  substance  is  a  mixture  of  liquid  oxygen  and 
liquid  nitrogen.     The  latter  boils  about   11°  lower  than  the 
former,  which  boils  at  —183°.     Therefore  nitrogen  distils  off 
first  when  liquid  air  is  allowed  to  evaporate,  leaving  nearly  pure 
oxygen.    This  is  stored  under  pressure  in  steel  tanks  and  brought 
on  the  market. 

310.  Erin's  Process.  —  Erin's  process,  formerly  used  tech- 
nically, is  a  method  of  obtaining  oxygen  from  the  air  by  means  of 
barium  oxide,  BaO.    This  oxide  unites  with  more  oxygen  at  a 
red  heat,  forming  barium  peroxide  : 


This  is  a  reversible  reaction,  which  at  a  constant  temperature 
will  go  in  one  direction  or  the  other  with  change  of  pressure. 
In  practice,  air  is  pumped  under  pressure  into  a  vessel  contain- 
ing BaO  at  700°,  and  when  all  the  oxide  has  been  changed  into 
Ba02  the  nitrogen  present  is  allowed  to  escape.  By  reducing 
the  pressure  with  a  vacuum  pump  the  BaO2  is  caused  to  decom- 
pose completely,  yielding  nearly  pure  oxygen. 

311.  Oxygen  from  Plants.  —  Growing  plants  absorb  carbon 
dioxide  from  the  air.  They  also  take  up  water  through  their 
roots.  In  some  manner,  not  fully  understood,  carbon  dioxide 
and  water  react  under  the  influence  of  sunlight  to  form  such 


1 86  Introduction  to  General  Chemistry 

principal  plant  constituents  as  starch,  cellulose,,  and  sugar, 
together  with  oxygen,  which  is  given  off  to  the  air.  The  per- 
centage of  oxygen  in  the  air  would  soon  decrease  if  it  were  not 
maintained  by  growing  plants. 

312.  The  Physical  Properties  of  Oxygen.— It  is,  of  course, 
obvious  that  oxygen  is  colorless,  odorless,  and  tasteless.    One 
liter  weighs  i.42Qg.  and  22.4  liters  about  32  g.,  corresponding 
to  the  formula  O2.    Liquid  oxygen  is  pale  blue  in  color;    it 
boils  at  —183°.    At  o°,  100  c.c.  of  water  dissolves  about  5  c.c. 
of  oxygen;  at  20°,  about  3  c.c.  (125). 

313.  The  Chemical  Properties  of  Oxygen. — We  have  already 
learned  that  combustion  was  first  explained  by  Lavoisier  in 
1774  as  due  to  union  of  the  burning  substance  with  the  oxygen 
of  the  air  (13-15).    All  the  elements  so  far  studied,  except 
fluorine,  form  oxides.    This  does  not  mean  that  all  these  ele- 
ments burn,  since  some  oxides,  like  those  of  chlorine  and  silver, 
can  only  be  made  indirectly  (172).    The  oxides  of  metallic 
elements,  by  union  with  water,  form  hydroxides  which  are 
bases,  for  example: 

CaO+H2O->Ca(OH)2. 

The  oxides  of  non-metallic  elements,  including  carbon,  sulfur, 
nitrogen,  phosphorus,  and  the  halogens  (except  flourine),  give, 
with  water,  acids.  The  following  equations  will  serve  as  illus- 
trations of  such  reactions,  some  of  which  have  already  been 
studied;  the  others  will  be  studied  later. 

COa+H20->H2CO3,  carbonic  acid, 

S03+H2O->H2SO4,  sulfuric  acid, 
N2Os-f-H2O->2HNO3,  nitric  acid, 
P2OS+3H2O->2H3PO4,  phosphoric  acid, 

I2OS+H2O->2HIO3,  iodic  acid. 

An  oxide  which  by  union  with  water  forms  an  acid  is  often 
called  the  anhydride  of  the  acid. 

314.  Respiration. — Animals  breathe  air  in  order  to  obtain 
oxygen.    The  blood  contains  a  complex  substance,  haemo- 
globin, which  forms  with  oxygen  a  compound,  oxyhaemoglobin, 
which  easily  decomposes  reversibly  into  oxygen  and  haemo- 


Hydrogen  and  Oxygen  187 

globin.  This  is  a  typical  equilibrium  reaction:  when  oxygen, 
at  the  pressure  at  which  it  exists  in  the  air,  comes  in  contact  with 
the  blood  in  the  lungs  the  compound  is  formed,  i  g.  of  haemo- 
globin uniting  with  1.3  c.c.  of  oxygen;  when  the  blood  reaches 
the  tissues,  which  take  up  oxygen,  the  compound  decomposes 
and  the  haemoglobin  is  carried  by  the  blood  back  to  the  lungs, 
where  it  again  takes  up  fresh  oxygen  from  the  air. 

315.  Uses  of  Oxygen.    The  Oxyacetylene  Torch. — The  use 
of  oxygen  in  the  oxyhydrogen  blowpipe  has  already  been  men- 
tioned.   By  substituting  acetylene  for  hydrogen  in  a  blowpipe 
similarly  constructed  we  get  the  oxyacetylene  torch,   which 
gives  an  intensely  hot  flame.     It  is  extensively  used  for  welding 
and  for  cutting  iron  and  steel. 

Oxygen  is  used  in  several  analytical  processes,  such  as  those 
studied  in  chapter  iv. 

Deposits  of  carbon  in  the  cylinders  of  gasoline  engines  are 
often  removed  by  burning  out  with  oxygen.  Since  iron  burns 
also  rather  readily  in  a  stream  of  oxygen,  care  must  be  taken  to 
avoid  injuring  the  cylinder  in  this  way. 

316.  Ozone. — When  a  silent  electric  discharge  passes  through 
oxygen  a  very  remarkable  change  is  produced;  there  is  a  decrease 
in  volume,  and  a  gas  having  a  powerful  irritating  odor  is  pro- 
duced.   The  new  gas  is  ozone.    The  simplest  form  of  apparatus 


Va 


s 


FIG.  43 

used  for  making  ozone  is  shown  in  Fig.  43.  It  is  a  double- walled 
glass  tube  having  the  outside  of  the  outer  tube  and  the  inside  of 
the  inner  tube  coated  with  tin  foil.  These  coatings  are  con- 
nected by  wires  to  the  terminals  of  an  induction  coil.  When  the 
coil  is  set  in  action  and  a  slow  stream  of  oxygen  is  passed  through 
the  space  between  the  outer  and  the  inner  tubes,  the  issuing  gas 
is  found  to  contain  ozone.  The  peculiar  odor  of  the  air  in  the 


1  88  Introduction  to  General  Chemistry 

neighborhood  of  powerful  electrical  machinery  is  due  to  ozone. 
Ozone  is  very  much  more  active  as  an  oxidizing  agent  than 
oxygen.  Mercury  shaken  with  ozone  is  very  quickly  oxidized. 
Ozone  also  sets  iodine  free  from  a  solution  of  an  iodide. 

If  nothing  but  oxygen  is  needed  to  produce  ozone  —  and  such 
is  actually  the  case  —  what  then  is  the  cause  of  the  remarkable 
change  in  properties?  In  the  first  place  it  was  noticed  that  a 
decrease  in  volume  occurs  when  ozone  is  formed  from  oxygen. 
On  the  other  hand,  when  ozone  is  changed  to  oxygen,  as  may  be 
done  by  heating  the  former,  the  volume  of  the  oxygen,  when  it 
is  again  cooled  to  the  original  temperature,  is  greater  than  that 
of  the  ozone.  In  fact,  three  volumes  of  oxygen  give  exactly  two 
of  ozone  and  vice  versa.  The  density  of  ozone  is  one-half 
greater  than  that  of  oxygen.  While  22.4  liters  of  oxygen  weigh 
32  g.,  the  same  volume  of  ozone  weighs  48  g.  Ozone  is  oxygen 
in  another  form.  If  for  oxygen  we  write  the  formula  O2,  we  must 
write  O3  as  the  formula  of  ozone.  The  molecules  of  ozone 
differ  from  those  of  oxygen  by  containing  three  instead  of  two 
atoms.  We  may  write  the  reversible  equation  for  the  relation 
between  oxygen  and  ozone  thus: 


When  ozone  acts  on  mercury,  for  example,  the  action  is  as  follows: 


Only  one-third  of  the  oxygen  of  ozone  is  active,  the  balance  changing 
into  ordinary  oxygen.  Iodides  are  oxidized  by  ozone,  thus: 

2KI+03+H20^2KOH+I2+02. 

The  liberated  iodine  may  be  recognized  by  its  action  on  starch. 
Minute  amounts  of  ozone  may  be  recognized  in  this  way, 
although  the  test  is  not  conclusive  proof  of  the  presence  of  ozone, 
since  many  other  substances  also  set  iodine  free  from  iodides. 

317.  Ozone  as  a  Germicide.  —  Since  ozone  is  a  very  powerful 
ozidizing  agent,  it  is  not  surprising  that  it  should  readily  destroy 
germs.  It  has  been  found  that  impure  water  containing  even 
1,000,000  bacteria  per  c.c.  is  completely  sterilized  by  intimate 


Hydrogen  and  Oxygen  189 

contact  with  an  equal  volume  of  air  containing  2  g.  of  ozone 
per  cubic  meter.  In  a  number  of  important  cities  of  Europe  the 
entire  municipal  water  supply  is  purified  by  means  of  ozone. 
Disinfection  by  chlorine  is  more  popular. 

318.  Hydrogen  Peroxide,  H2O2. — The  well-known  household 
antiseptic  and  disinfectant,  hydrogen  peroxide^  H202,  is  a  3  per 
cent    solution    of    this   substance   in   water.     Some   hydrogen 
peroxide  is  formed  by  the  action  of  sodium  peroxide  on  ice 
water,  thus: 

Na2O2+  2H2O  ±;  H2O2+  2NaOH. 

If  water  is  dropped  on  sodium  peroxide  the  material  becomes 
very  hot,  and  oxygen  and  water  instead  of  hydrogen  peroxide 
are  formed  (307).  This  is  because  the  latter  substance  easily 
decomposes  if  hot,  especially  in  the  presence  of  caustic  soda: 

2H2O2^2H2O-f-02. 

319.  Preparation  of  Hydrogen  Peroxide. — Barium  peroxide, 
Ba02,  the  formation  of  which  was  discussed  in  connection  with 
Erin's  process  of  making  oxygen  (310),  reacts  with  dilute  sulfuric 
acid  to  form  hydrogen  peroxide  and  barium  sulfate : 

BaO2+H2S04±;H2O2+BaS04. 

Since  barium  sulfate  is  insoluble  in  water  a  very  pure  solu- 
tion of  hydrogen  peroxide  is  easily  obtained.  The  reaction  is 
best  carried  out  by  adding  finely  powdered  barium  peroxide, 
suspended  in  water,  very  gradually  to  ice-cold,  diluted  sul- 
furic acid.  The  precipitate  of  barium  sulfate  is  allowed  to 
settle,  leaving  a  clear  solution  of  hydrogen  peroxide.  This 
solution  must  be  made  as  nearly  neutral  as  possible,  other- 
wise it  will  decompose  more  or  less  rapidly  into  oxygen  and 
water. 

By  cautious  evaporation,  at  a  moderate  temperature  in  a 
partial  vacuum,  a  dilute  solution  of  hydrogen  peroxide  may  be 
freed  from  most  of  its  water;  the  resulting  concentrated  solution 
when  cooled  to  — 10°  deposits  crystals  of  H2O2. 

320.  Properties  of  Hydrogen  Peroxide. — At  ordinary  tem- 
peratures pure  hydrogen  peroxide  is  a  colorless  liquid  which  will 


i  go  Introduction  to  General  Chemistry 

mix  with  water  in  all  proportions.  It  freezes  at  —  2°.  It  does 
not  boil  without  decomposition,  and  when  strongly  heated  it  is 
liable  to  explode,  water  and  oxygen  being  the  products.  The 
speed  of  decomposition  of  hydrogen  peroxide  at  ordinary 
temperatures  is  greatly  influenced  by  the  presence  of  other  sub- 
stances which  act  as  catalytic  agents.  Finely  divided  metals 
like  platinum  and  gold  cause  hydrogen  peroxide  to  decompose 
rapidly.  Manganese  dioxide  behaves  similarly.  In  these  reac- 
tions neither  the  metals  nor  the  manganese  dioxide  are  changed. 
They  are  catalytic  agents  (303). 

Hydrogen  peroxide  seems  to  have  the  property  of  an  acid, 
since  it  combines  with  some  bases  to  form  compounds  which 
may  be  considered  salts.  For  example,  with  barium  hydroxide, 
Ba(OH)2,  it  reacts  thus: 

H2O2+Ba(OH)2-j-6H2O±?BaO2-  8H2O. 

The  product  consists  of  white  crystals,  rather  difficultly  soluble 
in  water. 

Hydrogen  peroxide  is  often  used  as  a  bleaching  agent  for 
plant  and  animal  substances,  such  as  hair,  feathers,  silk,  ivory, 
and  straw. 

The  most  characteristic  property  of  hydrogen  peroxide  is  its 
great  tendency  to  give  up  oxygen  and  thus  to  act  on  substances 
capable  of  reacting  with  oxygen.  For  example,  with  hydriodic 
acid  it  gives  iodine  and  water: 

H202+2HI^I2+2H20. 

The  free  iodine  can  easily  be  recognized  by  the  blue  color  which 
it  gives  with  a  starch  solution.  Instead  of  using  hydriodic 
acid  we  may  use  a  solution  of  potassium  or  sodium  iodide  to 
which  hydrochloric  acid  has  been  added,  since  the  solution  will 
then  contain  some  hydriodic  acid  formed  as  follows: 

KI+HCl^KCl+HL 

321.  Detection  of  Hydrogen  Peroxide. — A  very  delicate 
reaction  which  serves  to  detect  small  quantities  of  hydrogen 


Hydrogen  and  Oxygen  191 

peroxide  is  that  which  occurs  when  a  solution  of  this  substance 
is  mixed  with  a  little  sulfuric  acid  and  a  very  dilute  solution  of 
potassium  dichromate,  K2Cr207.  The  latter  substance  contains 
the  element  chromium,  Cr,  as  one  of  its  constituents.  The 
solution  turns  blue,  and  when  a  little  ether  is  added  and  shaken 
up  with  the  blue  solution  the  ether  dissolves  the  blue  substance. 
If  the  mixture  is  allowed  to  stand  a  minute  or  two  the  blue  ether 
solution  separates  from  the  water  solution,  on  which  it  floats  as  a 
blue  layer. 

Other  reactions  of  hydrogen  peroxide  are  discussed  in  the 
following  chapter  (347,  348). 

322.  Peroxides  and  Dioxides. — We  have  just  learned  that 
hydrogen  peroxide  is  formed  by  the  action  of  dilute  acids  on 
Na202  and  Ba02,  and  we  might  therefore  be  inclined  to  expect 
that  we  should  also  get  H2O2  by  the  action  of  acids  on  Pb03  and 
Mn02.    But  this  is  not  the  case;   no  H2Oa  can  be  obtained  in 
any  way  from  these  last-mentioned  oxides.    For  this  reason 
these  oxides  of  lead  and  manganese  are  called  dioxides  to  dis- 
tinguish them  as  a  class  from  those  which  yield  H2O2  and  which 
are  called  peroxides.     Thus  we  call  BaO2  barium  peroxide  and 
Pb02  lead  dioxide. 

323.  Graphic  Formulae. — We  have  so  far  considered  that 
oxygen  has  a  valence  (183)  of  two,  or  is  bivalent,  since  in  water 
two  symbol  weights  of  hydrogen  are  united  with  one  of  oxygen. 
But  what  then  is  the  valence  of  oxygen  in  H2O2?    In  order  to  be 
able  to  answer  this  question,  we  must  consider  the  matter  of 
valence  from  the  standpoint  of  the  atomic-molecular  hypothesis. 
We  have  learned  (221)  that  the  molecule  of  water  is  made  up 
of  two  atoms  of  hydrogen  and  one  of  oxygen.    Since  all  the 
molecules  of  water  are  made  up  in  just  this  fashion,  it  would  seem 
to  follow  that  the  three  atoms  must  be  related  to  one  another 
in  some  very  definite  way.    We  may  think  of  them  as  being 
joined  to  one  another,  in  which  case  there  are  the  two  possibilities 
indicated  by  the  following  graphic  formulae  in  which  the  lines 
joining  the  symbols  are  called  bonds. 

(i)  H-H-0  (2)  H-O-H. 


Introduction  to  General  Chemistry 

The  first  graphic  formula  indicates  that  one  of  the  hydrogen 
atoms  is  attached  on  the  one  hand  to  the  atom  of  oxygen  and 
on  the  other  to  the  second  atom  of  hydrogen.  Formula  (2)  indi- 
cates that  it  is  the  oxygen  atom  which  is  attached  on  either  hand 
to  an  atom  hydrogen.  It  is  obvious  that  the  second  formula 
is  the  more  consistent,  since  in  it  both  atoms  of  hydrogen  are 
attached  by  single  bonds  to  the  atom  of  oxygen  which  holds  an 
atom  of  hydrogen  by  each  of  its  two  bonds.  In  formula  (i)  the 
middle  hydrogen  atom  is  represented  as  having  two  bonds,  while 
the  other  hydrogen  atom  and  also  the  atom  of  oxygen  are  shown 
as  having  but  one  bond  each.  Since  we  think  of  all  atoms  of 
hydrogen  as  being  alike,  we  must  reject  the  first  formula  in  favor 
of  the  second. 

When  viewed  in  the  above-mentioned  manner  the  valence  of 
an  element  is  seen  to  be  the  holding  capacity  of  its  atoms  for  atoms 
of  hydrogen  or  other  univalent  elements  like  chlorine.  We  may 
therefore  think  of  an  atom  of  oxygen  which  is  bivalent  as  having 
two  valence  bonds,  each  of  which  can  hold  one  atom  of  a  univalent 
element. 

324.  The  Graphic  Formulae  of  Peroxides. — We  are  now 
prepared  to  consider  the  question  of  the  valence  of  oxygen  in 
H2O2  and  the  graphic  formula  of  this  substance.  At  the  outset 
it  may  be  stated  that  the  valence  of  hydrogen  is  considered  by 
chemists  to  be  invariably  one.  If  the  valence  of  oxygen  is  taken 
to  be  two,  there  is  but  one  possible  graphic  formula,  namely: 

H-O-O-H. 

This  is  the  commonly  accepted  formula.  For  sodium  peroxide 
we  then  have  the  formula, 

Na-O-O-Na, 
and  for  barium  peroxide, 

B; 


Hydrogen  and  Oxygen  193 

In  their  dioxides,  manganese  and  lead  have  without  doubt 
a  valence  of  four,  or  are  tetravalent,  and  oxygen  is,  as  usual, 
bivalent.  We  therefore  write  for  these  oxides  the  formulae 

O=Mn=O  and  O=Pb  =  O, 

thereby  indicating  that  each  oxygen  atom  is  attached  to  an 
atom  of  manganese  or  lead  by  two  bonds,  or  in  other  words  by  a 
double  bond. 

The  monoxides  of  manganese,  MnO,  and  lead,  PbO,  have 
their  atoms  doubly  bound,  thus : 

Mn=OandPb=0. 
In  these  oxides  both  metals  are  bivalent. 


CHAPTER  XV 
OXIDATION  AND  REDUCTION 

325.  Oxidation.  —  When  a  substance  unites  with  oxygen  «it 
is  said  to  be  oxidized.  Hydrogen  when  burned  is  oxidized,  giv- 
ing water.  Metals  are  said  to  be  oxidized  when  they  combine 
with  oxygen;  for  example, 

2Cu+O2-»2CuO. 

By  certain  indirect  methods  a  low.er  oxide  of  copper,  cuprous 
oxide,  Cu2O,  can  be  made.  This  oxide  can  unite  with  more 
oxygen  if  heated  in  air  or  in  oxygen  and  form  the  common  oxide, 
CuO,  which  is  known  also  as  cupric  oxide,  in  order  to  distinguish 
it  from  the  lower  oxide: 


We  say  in  this  case  that  cuprous  oxide  has  been  oxidized  to  cupric 
oxide. 

The  action  of  oxygen  gas  on  hydrogen  chloride  at  a  high 
temperature  (239)  proceeds  according  the  equation 


and  in  consequence  we  say  that  the  hydrogen  chloride  has  been 
oxidized. 

326.  Oxidizing  Agents.  —  Very  often  substances  may  be 
oxidized  by  compounds  of  oxygen  as  well  as  by  oxygen  itself. 
For  example,  heated  copper  oxide,  CuO,  oxidizes  hydrogen  and 
all  its  compounds,  such  as  ammonia,  acetylene,  etc.  We  say, 
therefore,  that  copper  oxide  is  an  oxidizing  agent.  Any  sub- 
stance which  oxidizes  another  is  called  an  oxidizing  agent.  Lead 
dioxide  and  manganese  dioxide  are  powerful  oxidizing  agents, 
as  shown  by  the  fact  that  each  is  able  to  oxidize  hydrochloric 

acid, 

Mn03+4HCl->MnCL4-  Cl,+  2H2O.  (234) 

194 


Oxidation  and  Reduction  195 

Potassium  permanganate,  KMn04,  which  also  easily  oxi- 
dizes hydrochloric  acid,  is  one  of  the  most  powerful  of  all  oxidiz- 
ing agents.  It  acts  according  to  the  equation 

2KMn04+i6HCl->2KCl+2MnCl3+sCl2+8H20.        (235) 

Nitric  acid  and  its  salts  (104),  the  nitrates,  are  good  oxidiz- 
ing agents.  Gunpowder  is  a  mixture  of  finely  powdered  potas- 
sium nitrate,  charcoal,  and  sulfur.  The  explosion  of  gunpowder 
is  due  to  the  extremely  rapid  oxidation  of  the  charcoal  (carbon) 
and  sulfur  to  carbon  dioxide,  CO2,  and  sulfur  dioxide,  SO2-,  the 
oxygen  being  furnished  by  the  potassium  nitrate,  KNO3.  The 
other  products  of  the  explosion  are  nitrogen  and  potassium  sul- 
fide,  K2S. 

Potassium  chlorate,  KC103,  is  a  powerful  oxidizing  agent, 
readily  giving  up  all  its  oxygen  to  oxidizable  substances  and 
leaving  the  chloride,  KC1. 

Sulfuric  acid  (93)  is  capable  of  oxidizing  some  substances; 
hot  sulfuric  acid  acts  on  charcoal  thus: 

2H2S04+ C->C02+  2SO2+  2H2O. 

It  is  probable  that  each  molecule  of  sulfuric  acid  first  loses  one 
atom  of  oxygen  giving  a  molecule  of  sulfurous  acid,  H2S03,  and 
that  this  substance,  which  is  unstable,  then  decomposes  into 
sulfur  dioxide  and  water: 

H2SO3->SO2+H2O.  (286) 

327.  Reduction  and  Reducing  Agents. — When  hydrogen  is 
oxidized  by  hot  copper  oxide,  thus: 

CuO+H2->Cu-fH2O, 

the  copper  oxide  is  said  to  be  reduced  to  metallic  copper.  In 
consequence  we  call  hydrogen  a  reducing  agent.  Any  substance, 
as,  for  example,  acetylene  or  methane,  which  can  reduce  copper 
oxide  is  also  called  a  reducing  agent.  In  any  reaction,  if  one 
substance  is  oxidized  the  oxidizing  agent  is  by  necessity  reduced; 
oxidation  and  reduction  always  go  on  together.  All  substances 
which  are  acted  upon  by  oxidizing  agents  are,  of  course,  redu- 
cing agents. 


196  Introduction  to  General  Chemistry 

328.  Carbon  as  a  Reducing  Agent.  —  Since  charcoal,  which 
is  nearly  pure  carbon,  burns  readily,  it  is  capable  of  taking  up 
oxygen  from  oxidizing  agents  and  is  therefore  a  good  reducing 
agent.     A  mixture  of  powdered  copper  oxide  and  charcoal  reacts 
vigorously,  if  strongly  heated,  giving  copper  and  carbon  dioxide  : 

2CuO+C->2Cu+CO2. 

In  this  reaction  the  copper  oxide  is  the  oxidizing  agent  and  the 
charcoal  (carbon)  the  reducing  agent. 

Many  other  metallic  oxides  can  be  reduced  in  a  similar  man- 
ner by  carbon.  In  place  of  charcoal,  coke  or  coal,  which  are 
largely  carbon,  may  be  used.  Thus  ferric  oxide,  Fe203,  which 
in  the  form  of  the  mineral  hematite  is  the  most  important  ore 
of  iron,  is  reduced  by  coke  at  a  white  heat  to  metallic  iron.  The 
result  may  be  represented  by  the  equation 

2Fe2O3-f-3C->4Fe+3CO2, 

although  it  is  very  probable  that  the  reaction  is  less  simple  under 
the  conditions  actually  met  with  in  practice. 

329.  Carbon  Monoxide  as  a  Reducing  Agent.  —  When  carbon 
is  burned  in  a  deficient  supply  of  air,  carbon  monoxide,  CO,  is 
formed  instead  of  dioxide: 


This  is  a  colorless,  odorless,  and  very  poisonous  gas  which  will 
burn  with  a  nearlv  non-luminous  flame  to  form  carbon  dioxide, 

2CO+O2->2CO2. 

Some  oxidizing  agents  are  able  to  oxidize  carbon  only  to 
monoxide  and  not  to  dioxide.     Zinc  oxide  behaves  in  this  way: 

ZnO+C->Zn+CO. 

This  is  the  reaction  by  which  zinc  is  made  from  its  ores.     The 
reaction  between  ferric  oxide  and  carbon  can  also  give  carbon 

monoxide, 

Fe2O3+3C->2Fe+3CO. 

But  carbon  monoxide  can  also  reduce  ferric  oxide: 


Oxidation  and  Reduction  197 

The  last  two  equations  doubtless  represent  the  steps  by  which 
ferric  oxide  and  carbon  react  to  give  iron  and  carbon  dioxide 

(328). 

330.  Aluminum  as  a  Reducing  Agent.  —  Metallic  aluminum 
unites  vigorously  with  oxygen  at  a  white  heat,  although  it  has 
no  tendency  to  oxidize  in  the  air  at  ordinary  temperatures.  The 
burning  of  aluminum  occurs  thus  : 


When  a  mixture  of  powdered  aluminum  and  ferric  oxide  is 
strongly  heated  a  very  violent  reaction  takes  place,  giving  iron 
and  aluminum  oxide: 

2Al-hFe2O3->2Fe+Al2O3. 

The  mixture  of  aluminum  and  ferric  oxide  has  been  given  the 
trade  name  of  thermite  by  its  inventor,  Goldschmidt,  who  uses 
it  to  make  small  quantities  of  molten  iron  for  the  repair  of  broken 
iron  castings,  etc. 

Many  other  metallic  oxides  can,  also  be  reduced  by  aluminum. 

331.  Oxidation  Considered  as  a  Change  of  Valence.  —  We 
have  already  learned  (173)  that  iron  forms  two  series  of  com- 
pounds, ferrous  and  ferric,  as  illustrated  by  the  following  formulae: 

Ferrous  Compounds  Ferric  Compounds 

FeO  Fe2O3 

Fe(OH)2  Fe(OH)3 

FeCl2  FeCl3 

FeBr2  FeBr3 

Fe(N03)2  Fe(N03)3 

FeSO4  Fe2(SO4)3 

The  valence  of  iron  is  two  in  ferrous  compounds  and  three  in 
ferric.  According  to  this  usage  of  the  term  valence  we  should 
be  forced  to  say  that  the  valence  of  free  or  uncombined  iron  is 
zero. 

If  free  iron  is  changed  into  ferrous  oxide, 

2Fe-fO3->2FeO, 


198  Introduction  to  General  Chemistry 

it  is  oxidized,  and  its  valence  is  increased  from  zero  to  two. 
Moreover,  if  ferrous  oxide  is  changed  into  ferric  oxide, 

4FeO+Oa-»2Fe2O3, 

it  is  also  plain  that  the  iron  is  further  oxidized,  and  that  its 
valence  has  increased  from  two  to  three.  It  is  customary  to 
say  that  ferric  oxide  is  a  higher  oxide  of  iron  than  ferrous  oxide ; 
or  that  iron  in  ferric  oxide  is  in  a  higher  state  of  oxidation  than 
in  ferrous  oxide. 

In  the  case  of  iron  and  its  oxides  we  see  that  the  oxidation  of 
iron  and  the  increase  in  its  valence  go  hand  in  hand. 

With  respect  to  other  elements  that  unite  with  oxygen  we 
also  find  that  their  oxidation  results  in  an  increase  in  their 
valence.  A  few  additional  examples  will  help  to  illustrate  this 
point.  In  the  change  of  copper  into  cuprous  oxide,  Cu2O  (325), 
the  oxidation  of  the  copper  is  accompanied  by  an  increase  cf  its 
valence  from  zero  to  one.  In  cuprous  oxide  copper  is  univalent 
(146).  In  the  oxidation  of  cuprous  oxide  to  cupric  oxide, 

2Cu20-f-O2->4CuO,  (325) 

the  valence  of  copper  is  increased  from  one  to  two.  In  cupric 
oxide  copper  is  bivalent  (146). 

When  carbon  is  oxidized  to  carbon  monoxide, 

2C+02->2CO,  (329) 

the  valence  of  carbon  is  increased  from  zero  to  two  (carbon  is 
bivalent  in  carbon  monoxide).  In  the  oxidation  of  carbon 
monoxide  to  carbon  dioxide, 

2CO+O2-»2CO2,  (329) 

the  valence  of  carbon  is  increased  to  four,  carbon  becoming 
quadrivalent. 

332.  A  Broader  Meaning  of  the  Term  Oxidation. — Since  in 
the  change  of  any  ferrous  compound  into  the  corresponding 
ferric  compound  (173,  331)  the  valence  of  iron  always  increases 
from  two  to  three,  all  such  changes  may  well  be  considered  to 
be  of  the  same  class.  It  has  become  the  custom  among  chemists 


Oxidation  and  Reduction  199 

to  call  such  increase  of  valence  of  iron  an  oxidation  of  the  iron 
irrespective  of  the  nature  of  the  element  or  radical  combined 
with  the  iron.  Thus  in  the  reaction 

2FeCla+Cl3->2FeCl3, 

whereby  ferrous  chloride  is  changed  to  ferric  chloride,  we  say 
that  the  iron  has  been  oxidized.  The  only  question  that  should 
arise  here  is:  Why  call  this  increase  in  valence  of  iron  an  oxida- 
tion in  cases  where  no  oxygen  is  involved?  We  can  only  say 
that  it  is  a  custom  sanctioned  by  long  and  universal  usage. 

By  way  of  further  illustration  of  the  use  of  the  term  oxida- 
tion in  its  broader  sense  we  may  cite  the  following  examples. 
When  metallic  sodium  is  changed  into  chloride,  NaCl,  or  nitrate, 
NaNO3,  its  valence  is  increased  from  zero  to  one,  and  we  say 
that  the  sodium  has  been  oxidized.  When  zinc  is  changed  into 
oxide,  ZnO;  sulfate,ZnS04;  chloride,  ZnCl2;  or  nitrate,  Zn(N03)2 
(148),  we  say  that  the  zinc  has  undergone  oxidation;  and  further- 
more, since  in  all  these  compounds  zinc  is  bivalent,  we  say  that 
zinc  in  all  these  compounds  is  in  the  same  state  or  stage  of  oxida- 
tion. In  fact,  zinc  in  its  compounds  is  always  bivalent. 

333.  Review  of  Other  Elements  with  Variable  Valence. — 
Iron  is  not  the  only  element  having  a  variable  valence.  We 
have  already  seen  (179,  180)  that  mercury  also  forms  two  series 
of  compounds,  the  mercurous,  in  which  the  element  has  a  valence 
of  one,  and  the  mercuric,  where  the  valence  is  two,  as  illustrated 
by  the  following  formulae: 

Mercurous  Compounds  Mercuric  Compounds 

Hg30  HgO 

HgCl  HgCla 

Hgl  HgI2 

HgN03  Hg(N03)2 

Hg2S04  HgS04 

Mercurous  compounds  are  converted  into  mercuric  by  oxida- 
tion, and  mercuric  into  mercurous  by  reduction. 

Copper  also  forms  two  series  of  compounds,  cuprous  and 
cupric.  We  know  cuprous  oxide,  CuaO,  and  cuprous  chloride, 
CuCl,  as  well  as  the  commoner  cupric  compounds,  such  as  cupric 


2oo  Introduction  to  General  Chemistry 

oxide,  CuO,  cupric  chloride,  CuCl2,  cupric  sulfate,  CuSO4, 
etc.  (165). 

It  will  be  noted  that  compounds  representing  the  lower  state 
of  oxidation  have  names  ending  in  ous,  while  those  corresponding 
to  the  higher  state  of  oxidation  end  in  ic. 

334.  Another    Class    of    Oxidation    Reactions. — We    have 
given  as  one  illustration  of  an  oxidation  reaction  the  action  of 
manganese  dioxide  on  hydrochloric  acid: 

MnO2+4HCl->MnCl2+Cl2+2H2O  (326) 

In  this  reaction  the  manganese  dioxide  is  the  oxidizing  agent. 
The  chlorine  of  the  hydrochloric  acid  in  being  set  free  has  its 
valence  decreased  from  one  to  zero.  The  valence  of  the  hydro- 
gen remains  unchanged  in  the  reaction:  it  is  neither  oxidized 
nor  reduced.  Now  we  see  that  when  combined  chlorine,  as  in  a 
chloride,  is  set  free  its  valence  is  decreased,  and  that  in  this  change 
the  chlorine  is  oxidized.  In  the  case  of  metallic  elements  oxida- 
tion involves  increase  of  valence  (331).  We  see,  therefore,  that 
in  the  case  of  chlorine,  a  non-metallic  element,  its  oxidation  in 
cases  like  that  cited  involves  a  decrease  in  its  valence.  Other 
non-metallic  elements,  like  bromine,  iodine,  and  sulfur,  when  set 
free  by  oxidizing  agents  from  their  compounds  with  hydrogen 
or  metals  are  oxidized,  while  at  the  same  time  there  occurs  a 
decrease  in  valence. 

335.  Reduction   and    Change    of    Valence. — That   valence 
changes  accompany  reduction,  as  well  as  oxidation,  will  be  at 
once  apparent  by  the  consideration  of  any  reaction  in  which 
reduction  takes  place.     Take,  for  example,  the  simple  case  of 
the  reduction  of  cupric  oxide  by  hydrogen, 

CuO+H2->Cu-{-H2O.  (327) 

The  copper  is  reduced  to  the  free  state,  its  valence  changing 
from  two  to  zero,  while  the  oxygen  merely  changes  partners 
without  change  of  valence  and  is  therefore  neither  oxidized  nor 
reduced.  When  any  oxide  of  a  metal  is  reduced  by  hydrogen, 
carbon,  or  any  other  reducing  agent,  the  valence  of  the  metal  is  also 
reduced  or  lowered. 


Oxidation  and  Reduction  201 

On  the  other  hand,  in  the  reaction 
C12+H2-»2HC1 

we  see  that  the  reduction  of  the  non-metallic  element  chlorine 
is  accompanied  by  an  increase  in  its  valence  from  zero  to  one. 
In  general,  when  a  non-metallic  element  like  a  halogen,  sulfur,  or 
oxygen  itself  unites  with  hydrogen  or  a  metal,  the  non-metal  is 
reduced,  while  concurrently  its  valence  is  increased. 

336.  Classification  of  Valence  Changes. — In  analyzing  the 
foregoing  cases  we  have  seen  that  there  is  a  difference  in 
the  behavior  of  metals  and  non-metals.  If  we  consider  the 
many  compounds  which  we  have  studied,  we  shall  see  that 
metals  and  hydrogen  unite  with  non-metals  more  or  less 
readily  but  do  not  unite  with  each  other,  at  least  in  the  cases 
studied.  On  the  other  hand  non-metals  not  only  unite  with 
metals  but  also  unite  among  themselves,  as  in  the  case  of 
carbon  and  oxygen  or  sulfur  and  oxygen  (286). 

In  all  but  the  rarest  cases  the  valence  of  an  atom  in  com- 
bination may  be  found  directly  or  indirectly  from  the  following 
rules,  which  are  established  by  experience:  hydrogen  (sodium 
and  potassium)  are  univalent  in  all  their  compounds;  oxygen 
is  bivalent  in  all  its  compounds  except  the  peroxides.  The 
latter  have  definite  characteristics  and  may  be  identified.  The 
classification  of  the  changes  of  valence  may  be  made  systemati- 
cally on  the  basis  of  the  change  of  the  valence  of  an  atom  toward 
a  metal  or  non-metal  as  indicated  in  the  table  below.  Whether 
the  atom  in  question  is  a  metal  or  a  non-metal  does  not  matter. 


For  the  Atom  in 
Question 

In  Oxidation 

In  Reduction 

Valence  toward 
non-metals.  .  . 
Valence  toward 
metals  or  hyr- 
drogen  

Increases 
Decreases 

Decreases 
Increases 

If  the  reactions  of  reagents  are  classified  according  to  this  scheme 
those  which,  are  found  to  undergo  oxidation  or  reduction,  and 


2O2  Introduction  to  General  Chemistry 

so  are  themselves  either  reducing  agents  or  oxidizing  agents, 
will  be  found  actually  to  have  similar  chemical  activities. 

337.  Two  Important  Kinds  of  Reactions. — But  very  few  of 
the  reactions  of  substances  in  solution  studied  in  chapters  prior 
to    the   present   involve    oxidation    and    reduction.     In    such 

reactions  as 

HCl+NaOH->NaCl-hH2O, 
AgN03+HCl->AgCl+HN03,  (169) 

and 

FeCl3+3NaOH-»Fe(OH)3+3NaCl,  (173) 

neither  oxidation  nor  reduction  occurs.  These  are  called  double- 
decomposition  reactions.  In  such  reactions  no  element  changes 
its  valence. 

If  oxidation  and  reduction  take  place,  the  reaction  is  of  a 
distinctly  different  kind.  In  such  reactions  two  or  more  elements 
change  valence. 

338.  Intensity  of  Activity  of  Oxidizing  and  Reducing  Agents. 
— Both  oxidizing  and  reducing  agents  differ  greatly  in  their 
intensity  of  activity.    For  example,  manganese  dioxide  will 
oxidize  cold,  dilute  hydrochloric  acid,  but  oxygen  gas  will  not. 
In  consequence  we  say  that  manganese  dioxide  is  a  stronger 
or  more  powerful  oxidizing  agent  than  oxygen  itself.    When  we 
find,  as  we  may  readily  do  by  experiment,  that  dilute  nitric  acid 
will  oxidize  a  ferrous  salt  to  a  ferric  salt  and  that  dilute  sulfuric 
acid  will  not  do  so,  we  conclude  that  nitric  acid  is  a  stronger  or 
better  oxidizing  agent  than  sulfuric  acid.    When  we  know  that 
hydriodic  acid  will  reduce  sulfuric  acid,  but  that  hydrochloric 
acid  will  not  do  so,  we  conclude  that  hydriodic  acid  is  a  better 
reducing  agent  than  hydrochloric  acid. 

339.  Hydrogen  Sulfide,  H2S. — Iron  and  sulfur  unite  directly 
at  a  red  heat  to  form  ferrous  sulfide,  FeS : 

Fe+S->FeS. 

This  is  a  black  solid,  which  is  insoluble  in  water.  It  reacts 
readily  with  hydrochloric  acid  to  give  ferrous  chloride  and 
hydrogen  sulfide: 

FeS-f-  2HCl->FeCla+H2S. 


Oxidation  and  Reduction 


203 


Hydrogen  sulfide  is  a  colorless  gas,  of  which  water  dissolves 
three  to  four  times  its  own  volume.  It  has  a  very  disagreeable 
odor,  resembling  rotten  eggs,  and  is  extremely  poisonous. 
Fatal  accidents  have  often  occurred  from  breathing  the  gas. 
Hydrogen  sulfide  is  a  very  powerful  reducing  agent.  In  water 
solution  it  is  easily  oxidized  by  atmospheric  oxygen,  giving  sulfur 
and  water: 

2H2S+O2-»2S+2H2O. 

A  water  solution  of  hydrogen  sulfide  reacts  rapidly  with  iodine 
to  form  hydriodic  acid  (265)  and  sulfur: 

H2S+I2->2HI+S. 

This  reaction  furnishes  a  very  good  practical  method  for  making 
hydriodic  acid.  We  have  only  to  pass  hydrogen  sulfide  gas  into 
water  containing  powdered  iodine.  When  all  the  iodine  has 
been  reduced,  the  solid  sulfur  can  be  filtered  out,  giving  a 
clear,  colorless  filtrate  which  contains  only  hydriodic  acid  and 
water. 

Hydrogen  sulfide  readily  reduces  dilute  sulfuric  acid,  which 
is  but  a  very  mild  oxidizing  agent  capable  of  oxidizing  only  the 
most  active  reducing  agents;  the  products  are  sulfurous  acid, 
H2S03,  sulfur,  and  water: 

H2S-f-H2SO4->H2S03+S-fHaO. 

Hydrogen  sulfide  can  be  made  by  the  action  of  sulfuric  acid 
on  ferrous  sulfide,  thus: 

FeS+H2S04-»FeSO4+H2S ; 

but  this  is  not  advisable  in  practice  because  of  the  interaction  of 
the  sulfuric  acid  with  the  hydrogen  sulfide.  Hydrochloric  acid 
is  the  best  acid  to  use  in  making  hydrogen  sulfide. 

Hydrogen  sulfide  is  oxidized  by  all  except  the  very  mildest 
oxidizing  agents.  As  a  final  example  of  its  behavior,  its  action 
on  ferric  salts  may  be  given.  These  are  reduced  to  ferrous  salts, 

thus: 

2FeCl3+H2S->2FeCl2+S+2HCl. 


204  Introduction  to  General  Chemistry 

340.  Sulfurous  Acid,  H2SO3.— When  sulfur  burns,  it  forms 
sulfur  dioxide,  SO2,  a  colorless  gas,  with  a  strong  odor: 

S+O2-»SO2.  (286) 

Sulfur  dioxide  is  very  soluble  in  water,  with  which  it  unites 
partially  to  form  sulfurous  acid: 

SO2+H2O^H2SO3.  (286) 

This  reaction  is  reversible;  by  boiling  a  solution  of  sulfurous 
acid  the  latter  can  be  completely  decomposed  and  all  sulfur 
dioxide  driven  off. 

Sulfurous  acid  is  a  reducing  agent  which,  when  oxidized,  is 
converted  into  sulfuric  acid.  It  reacts  slowly  with  atmospheric 
oxygen,  thus: 

2H2SO3+O2->2H2SO4. 

It  is  rapidly  oxidized  by  chlorine,  which  is  thereby  reduced 
to  hydrochloric  acid: 

H2S03+C12+H2O>H2SO4+2HC1. 

Manganese  dioxide  and  sulfurous  acid  react  as  follows : 
MnO2+H2SO3->MnSO4+H2O. 

Ferric  salts  are  reduced  to  ferrous  salts  by  sulfurous  acid,  as 
illustrated  by  the  following  equation: 

Fe2(S04)3+H2SO3+H2O->2FeS04+2H2SO4. 

341.  Hydrochloric,  Hydrobromic,  arid  Hydriodic  Acids  as 
Reducing  Agents. — These  acids  in  water  solution  can  all  be 
oxidized,  and  are  therefore  to  be  considered  as  reducing  agents. 
Hydriodic  acid  is  the  most  easily  oxidized  of  the  three,  and  is 
therefore  the  best  or  most  powerful  reducing  agent.     It  is 
oxidized  by  atmospheric  oxygen  (265),  which  has  no  action 
whatever  on  a  water  solution  of  hydrochloric  acid.    The  latter 
substance  acts  as  a  reducing  agent  only  with  respect  to  the  most 
powerful  oxidizing  agents,  such  as  manganese  (234)  and  lead 
dioxides   (235)   and  potassium  permanganate   (236).     Hydro- 


Oxidation  and  Reduction  205 

bromic  acid  is  a  better  reducing  agent  than  hydrochloric  acid, 
but  not  as  powerful  as  hydriodic  acid. 

Hydrobromic  acid,  as  a  reducing  agent,  reacts  with  concen- 
trated sulfuric  acid,  as  an  oxidizing  agent,  as  follows, 

2HBr+H2SO4->Br2+SO2+2H2O, 

forming  free  bromine,  sulfur  dioxide,  and  water. 

Hydriodic  acid  reacts  even  more  vigorously  with  concentrated 
sulfuric  acid.  In  this  case  the  products  vary  according  to  the 
proportions  taken,  but  the  reduction  of  the  sulfuric  acid  may  go 
as  far  as  the  formation  of  free  sulfur  and  hydrogen  sulfide.  The 
possible  reactions  are  represented  in  the  following  equations: 

H2SO4+  2ffl-»I2+SO2+  2H2O, 

H2S04+6HI->3I2+S+4H20, 

H2S04+8HI->4I2+H2S+4H20. 

It  will  now  be  understood  why  roundabout  methods  are  used 
to  prepare  hydrogen  bromide  and  hydrogen  iodide  instead  of 
the  simple  reaction  with  concentrated  sulfuric  acid  and  a  salt, 
as  is  done  in  the  preparation  of  hydrogen  chloride. 

342.  Manganese  and  Its  Compounds. — Manganese,  Mn,  is  a 
metallic  element  which  in  the  free  form  resembles  iron  rather 
closely.  Its  principal  ore  is  the  dioxide,  Mn02,  called  by 
mineralogists  pyrolusite.  Manganese  forms  a  series  of  salts 
corresponding  to  the  salts  of  ferrous  iron;  among  such  we  have 
manganous  chloride,  MnCl2,  manganous  nitrate,  Mn(N03)2,  and 
manganous  sulfate,  MnS04.  These  salts  are  pale  pink  in  color 
and  are  easily  soluble  in  water.  The  dilute  solutions,  which  are 
almost  colorless,  give  with  sodium  hydroxide  white  precipitates 
of  manganous  hydroxide : 

MnCl2+2NaOH-^Mn(OH)2+2NaCl. 

This  hydroxide  corresponds  to  an  oxide,  MnO: 
Mn(OH)2->MnO+H2O. 

In  all  these  compounds  except  the  dioxide,  MnO2,  manganese  is 
bivalent;  in  the  dioxide  it  is  quadrivalent. 


206  Introduction  to  General  Chemistry 

Manganese  forms  a  variety  of  compounds  of  a  very  different 
character  from  the  ones  just  mentioned;  of  these  the  most 
important  is  potassium  permanganate. 

343.  Potassium  Permanganate,  KMnO4. — This  substance  is 
the  potassium  salt  of  permanganic  acid,  HMnO4,  in  which 
manganese  acts  as  an  acid-forming  element.    The  salt  is  made 
from  manganese  dioxide  and  potassium  hydroxide  by  compli- 
cated reactions  which  need  not  be  considered  at  present.    It 
forms  dark-purple  crystals  which  dissolve  in  water  to  form  a 
purple  solution  having  nearly  the  color  of  the  vapor  of  iodine. 
It  is  a  very  important  substance  and  is  one  of  the  most  powerful 
of  all  oxidizing  agents. 

We  have  already  learned  that  potassium  permanganate 
oxidizes  hydrochloric  acid,  thus: 

2KMn04+  i6HCl-»2KCl+  2MnCl2+  5C12+ 8H2O.        (235) 

It  can  also  oxidize  almost  any  substance  which  is  capable  of 
being  oxidized  in  solution.  Two  additional  examples  may  be 
given  as  illustrations: 

2KMnO4+5H2S03->K2SO4-|-2MnSO4-f2H2SO4+3H2O, 
2KMn04+ioFeSO4+8H2SO4->K2SO4+2MnSO4+sFe2(S04)3+8H20. 

In  the  last  reaction  the  sulfuric  acid  acts  neither  as  a  reducing 
nor  an  oxidizing  agent,  but  is  used  to  keep  the  solution  acid.  The 
sulfate  radical  is  not  decomposed  in  the  reaction. 

From  the  foregoing  equations  it  is  apparent  that  when  two 
molecules  of  permanganate  change  to  manganese  sulfate  or 
chloride,  they  change  the  valence  of  ten  atoms  (chlorine  in 
hydrochloric  acid  or  iron  in  a  ferrous  salt)  by  one  unit  of  valence 
each,  or  of  five  atoms  by  two  units  of  valence  each  (sulfur  in 
sulfurous  acid).  This  relationship  exists  because  the  valence 
of  manganese,  which  is  seven  in  permanganate,  changes  to  two 
in  manganese  sulfate  or  chloride. 

344.  Chromium  and  Its  Compounds. — The    element   chro- 
mium, Cr,  is  a  hard  metal,  resembling  iron  in  appearance.     It 
forms  a  series  of  salts  of  which  chromic  chloride,  CrCl3,  and 
chromic  sulfate,  Cr2(S04)3,  are  typical  examples.    Solutions  of 


Oxidation  and  Reduction  207 

chromic  salts  are  either  green  or  violet  in  color,  according  to  the 
method  of  preparation.  These  solutions  give  with  ammonium 
hydroxide  bluish  precipitates  of  chromic  hydroxide : 

CrCl3+3NH4OH->Cr(OH)3+3NH4Cl. 

The  hydroxide  when  strongly  heated  gives  chromic  oxide: 
2Cr(OH)3-»CrA+3H20. 

It  will  be 'seen  that  chromic  salts  are  analogous  to  ferric  salts 
and  that  in  these  compounds  chromium  is  trivalent. 

345.  Chromates  and  Bichromates. — When  chromic  oxide  is 
fused  with  sodium  nitrate  or  sodium  peroxide,  sodium  chromate, 
Na2Cr04,  is  formed.    This  is  a  bright-yellow  crystalline  salt, 
readily  soluble  in  water.     It  may  be  considered  as  derived  from 
chromic  acid,  H2Cr04.    Potassium  chromate,  K2Cr04,  is  also  a 
yellow  crystalline  salt  which  is  readily  made  by  methods  similar 
to  those  that  give  the  sodium  salt. 

A  solution  of  potassium  chromate,  which  is  bright  yellow  in 
color,  turns  deep  orange  when  mixed  with  sulfuric  acid.  The 
solution  contains  potassium  dichromate,  K2Cr207,  which  has 
been  formed  thus: 

2K2Cr04+H2S04->K2Cr207-|-K2S04+H20. 

Potassium  dichromate  forms  orange-colored  crystals,  which 
dissolve  in  water  to  form  an  orange-colored  solution.  In  the 
foregoing  reaction  we  might  have  expected  to  get  potassium 
hydrogen  chromate,  KHCr04;  but  if  this  salt  is  first  formed  it 
decomposes  at  once,  as  follows: 

2KHCr04->K2Cr2O7+H2O. 

Sodium  dichromate,  Na2Cr207,  orange-colored  crystals,  can 
be  made  in  a  similar  manner  from  sodium  chromate. 

346.  Chromates  and  Bichromates  as  Oxidizing  Agents. — 
Solutions  of  either  chromates  or  dichromates  are  oxidizing 
agents.    More  commonly  a  strongly  acid  solution  is  used.     Thus 
with  hydrogen  sulfide  an  acid  solution  of  potassium  chromate 


208  Introduction  to  General  Chemistry 

(potassium  dichromate)  reacts  to  form  sulfur  and  chromium 
salts: 

2K2CrO4+3H2S+ioHCl->4KCl+2CrCl3+8H20+3S; 

or  if  the  equation  is  written  for  the  dichromate  we  have 
K2Cr207+3H2S+8HCl->2KCl+2CrCl3+7H2O+3S. 

With  sulfurous  acid  as  a  reducing  agent  the  reaction  yields 
sulf uric  acid  and  chromium  sulf ate : 

2K2Cr04+3H2S03+2H2SO4->2K2SO4+Cr2(SO4)3+5H2O. 

Apparently  two  molecules  of  potassium  chromate  (or  one  of  the 
dichromate)  can  cause  six  units  of  valence  change  on  other 
atoms,  three  of  sulfur  in  H2S  if  free  sulfur  is  the  product,  or 
three  of  sulfurous  acid  to  form  sulf  uric  acid. 

It  is  plain  that  permanganates  are  more  powerful  oxidizers 
than  chromates  or  dichromates,  since  the  first  can  oxidize 
hydrochloric  acid  and  the  second  cannot,  except  in  very  con- 
centrated acid  solution. 

347.  Hydrogen  Peroxide  as  an  Oxidizing  Agent. — We  have 
already  learned  (318)  that  hydrogen  peroxide  easily  decomposes 
into  water  and  oxygen,  and  that  for  this  reason  it  acts  as  an 
oxidizing  agent.  Its  action  on  hydriodic  acid  was  shown  to  take 

place  thus: 

H2O2+  2HI-»I2+  2H20.  (320) 

Sulfurous  acid  is  readily  oxidized  to  sulf  uric  acid : 
H2O2+H2S03-»H2SO4+H2O. 

Ferrous  salts  are  oxidized  to  ferric  salts  as  the  following 
equation  will  illustrate: 

H202+2FeS04+H2S04->Fe2(SO4)3+2H2O. 

Lead  forms  with  sulfur  lead  sulfide,  PbS,  a  black  substance, 
almost  insoluble  in  water.  It  is  obtained  as  a  black  precipitate 
by  the  action  of  hydrogen  sulfide  on  a  solution  of  a  lead  salt: 

Pb(NO3)2+H2S->PbS+2HNO3. 


Oxidation  and  Reduction  209 

Hydrogen  peroxide  oxidizes  lead  sulfide  to  lead  sulfate  (167)  : 
PbS+4H2O2->PbSO4+4H2O. 

Since  lead  sulfate  is  white,  the  effect  of  the  action  is  easily  seen. 
The  blackening  of  old  oil  paintings  is  due  to  the  gradual  conver- 
sion of  the  lead  compounds  that  have  served  as  ingredients  of 
the  paint  into  lead  sulfide  by  the  action  of  sulfur  compounds 
occurring  in  the  air.  Blackened  paintings  are  often  restored 
to  their  original  colors  by  treating  them  with  hydrogen  peroxide, 
which  converts  the  black  lead  sulfide  into  white  lead  sulfate. 

It  has  already  been  mentioned  that  animal  and  vegetable 
substances  are  bleached  by  hydrogen  peroxide.  The  exact 
nature  of  the  changes  that  occur  in  such  reactions  is  not  in  general 
known,  but  it  is  safe  to  conclude  that  they  are  processes  of  oxida- 
tion which  convert  colored  into  colorless  substances. 

348.  The  Reducing  Action  of  Hydrogen  Peroxide.  —  The 
action  of  hydrogen  peroxide  on  silver  oxide  yields  free  silver, 
and  we  may  say  that  the  silver  oxide  has  been  reduced. 

H2O2+  Ag2O-»  2Ag+H2O+O2. 

Another  important  reaction  of  this  class  is  found  in  the  action 
of  hydrogen  peroxide  on  potassium  permanganate  in  acid  solu- 
tion, which  takes  place  thus: 

5H2O2+  2KMnO4-f  3H2SO4->  K2SO4+  2MnSO4+8H2O+5O2. 


The  products  are  the  colorless  solution  of  the  sulfates  of  potas- 
sium and  manganese  in  addition  to  free  oxygen. 

349.  Hypochlorous  Acid,  HC1O.  —  It  is  probable  that  chlorine 
reacts  reversibly  with  water  in  which  it  is  dissolved  to  form 
hydrochloric  acid  and  hypochlorous  acid,  HC10,  thus: 

C12+H20±5HC1+HC10. 

Since  this  is  a  reversible  reaction,  all  four  substances  are  con- 
tained in  equilibrium  in  a  solution  of  chlorine  in  water.  Hypo- 
chlorous  acid  is  very  unstable,  that  is,  it  easily  decomposes, 
and  for  this  reason  it  cannot  be  obtained  except  in  the  form  of  a 
dilute  water  solution.  It  has  only  very  weak  acid  properties 


210  Introduction  to  General  Chemistry 

and  cannot  even  decompose  calcium  carbonate,  which  is  acted 
upon  by  almost  all  other  acids.  As  is  well  known,  hydrochloric 
acid  reacts  with  calcium  carbonate  as  follows: 

2HCl+CaCO3-»CaCl2+CO2+H2O.  (163) 

As  a  matter  of  fact,  when  calcium  carbonate  is  added  to  chlorine 
water  it  reacts  as  follows : 

CaC03-f-2Cl2+H20->  2HC10+CaCl2+C02. 

From  the  resulting  solution  hypochlorous  acid  mixed  with  much 
water  vapor  can  be  driven  off  by  cautious  heating;  the  condensed 
vapor  forms  a  dilute  solution  of  hypochlorous  acid.  This  reac- 
tion seems  to  prove  that  chlorine  and  water  react  to  form  hydro- 
chloric and  hypochlorous  acids. 

350.  Hypochlorites. — If  chlorine  is  passed  into  a  cold,  dilute 
solution  of  sodium  hydroxide,  sodium  chloride  and  sodium  hypo- 
chlorite,  NaCIO,  are  formed: 

C12+  2NaOH  -»NaCl+NaC10+H3O. 

This  is  exactly  what  we  should  expect  if  both  acids  which  result 
from  the  action  of  chlorine  on  water  are  neutralized  by  the 
sodium  hydroxide.  Chlorine  and  potassium  hydroxide  react 

similarly: 

C12+2KOH->KC1+KC10+H2O. 

351.  Bleaching  Powder. — The  action  of  chlorine  gas  on  solid 
slaked  lime,  calcium  hydroxide,  takes  place  thus: 

Cl2+Ca(OH)2->CaCl(OCl)+H2O. 

The  product  of  the  reaction  is  a  white  powder  known  as  chloride 
of  lime  or  bleaching  powder.  It  is  a  mixed  salt,  a  chloride  and 
hypochlorite  of  calcium.  It  is  extensively  used  in  the  bleaching 
of  cotton  goods  and  for  a  variety  of  other  purposes.  Before 
taking  up  the  chemical  behavior  of  hypochlorous  acid  and  hypo- 
chlorites  it  will  be  of  interest  to  consider  the  formation  of  these 
substances  from  the  standpoint  of  oxidation  and  reduction. 


Oxidation  and  Reduction  211 

352.  The  Oxidation  Products  of  Chlorine. — By  the  action  of 
chlorine  gas  on  dry  mercuric  oxide,  HgO,  chlorine  monoxide, 
C120,  a  brownish-yellow  gas,  is  obtained.     It  is  obvious  that  in 
this  reaction  the  chlorine  has  been  oxidized.    Now  this  oxide 
of  chlorine  unites  with  water  to  form  hypochlorous  acid, 

Cl20-fH2O-»2HC10. 

The  relation  between  chlorine  monoxide  and  hypochlorous 
acid  is  similar  to  that  between  sulfur  dioxide  and  sulfurous  acid: 

S02+H20-»H2S03. 

Chlorine  and  sulfur  also  show  similar  behavior  in  that  each  forms 
compounds  with  hydrogen  and  metals,  namely  chlorides  and 
sulfides. 

353.  The  Formation  of  Chlorates. — Hypochlorites  are  very 
unstable  salts.    A  warm,  concentrated  solution  of  sodium  hypo- 
chlorite  changes  more  or  less  rapidly  into  sodium  chloride  and 
sodium  chlorate,  NaClO3,  according  to  the  equation, 

3NaC10->  2NaCl-f-NaC103. 

Potassium  hypochlorite  changes  in  a  similar  fashion,  yielding 
potassium  chlorate,  KC103. 

Sodium  and  potassium  chlorates  are  powerful  oxidizing 
agents,  since  they  contain  large  proportions  of  easily  liberated 
oxygen.  When  the  dry  crystals  are  heated  they  decompose 
finally  into  chlorides  and  oxygen: 

2KC103->2KC1+302. 

This  reaction  takes  place  in  two  stages  (306).  The  first  change 
gives  rise  to  a  perchlorate,  KC1O4,  thus: 

ioKC103-»4KCl+6KC104-h302. 

Until  recently  potassium  chlorate  was  used  extensively,  and 
sodium  chlorate  was  rarely  seen.  The  reason  was  twofold:  in 
the  first  place  potassium  chlorate  was  made  very  largely  in 
Germany,  where  potassium  compounds  are  cheap  on  account 
of  the  immense  potash  deposits  found  in  that  country;  and  in 


212  Introduction  to  General  Chemistry 

the  second  place  sodium  chlorate,  being  more  soluble,  is  more 
difficultly  purified  than  the  potassium  salt.  Since  the  war  began 
there  has  been  a  shortage  of  potash,  because  no  other  country 
besides  Germany  has  much  easily  accessible  potash.  As  a 
consequence  the  manufacture  of  sodium  salts  has  been  stimu- 
lated, and  since  1915  there  has  been  an  abundant  supply  of 
sodium  chlorate.  This  can  be  used  advantageously  in  place  of 
potassium  chlorate  for  nearly  all  purposes. 

354.  Chloric  Acid  and  Chlorine  Dioxide. — Potassium  and 
sodium  chlorates  are  salts  of  chloric  acid,  HC103.    This  is  a  very 
unstable  acid,  which  is  known  only  in  dilute  solution.     Upon 
evaporation  of  the  solution  the  acid  decomposes,  giving  chlorine 
dioxide,  C102,  and  other  products. 

If  a  few  drops  of  concentrated  sulphuric  acid  are  poured  on 
a  small  crystal  of  sodium  chlorate  in  a  dry  test  tube,  a  yellow 
gas  forms,  which  explodes  with  violence  a  few  seconds  later. 
This  dangerous  experiment  should  be  performed  with  great 
caution.  The  yellow  gas  is  chlorine  dioxide,  C1O2,  which  was 
formed  by  the  decomposition  of  the  chloric  acid  set  free,  thus : 

NaClO3+H2SO4->NaHS04+HClO3. 

The  explosion  of  chlorine  dioxide  is  due  to  decomposition  into 
its  elements: 

2C1O2->C12+2O3. 

Chloric  acid  is  a  powerful  oxidizing  agent.  For  example,  it 
changes  lead  sulfide  to  lead  sulfate  (167).  This  operation  is 
usually  carried  out  by  adding  a  few  crystals  of  sodium  chlorate 
and  dilute  hydrochloric  acid  to  the  black  lead  sulfide.  The 
dark  color  is  seen  to  change  slowly  to  the  white  of  the  sulfate : 

3PbS+4HC103  ->  3PbS04+4HCl. 

355.  Perchlorates  and  Perchloric  Acid. — -Perchlorates   are 
formed  by  heating  .chlorates  gently  (306,  353). 

Sodium  perchlorate,  NaCl04,  and  potassium  perchlorate, 
KC104,  are  white  crystalline  salts.  They  decompose  completely 
into  chlorides  and  oxygen  at  dull-red  heat.  For  example, 

.NaClO4->NaCH-2(V 


Oxidation  and  Reduction  213 

Ammonium  perchlorate,  NH4C1O4,  is  made  by  neutralizing 
perchloric  acid,  HC104,  with  ammonia.  It  is  used  as  an  oxidizing 
agent  and  as  a  very  powerful  explosive. 

When  powdered  sodium  or  potassium  perchlorate  is  mixed 
with  concentrated  sulf  uric  acid  and  cautiously  heated  in  a  small 
retort  (104,  Fig.  24),  perchloric  acid,  HC104,  is  distilled  from  the 
mixture.  This  experiment  should  not  be  made  by  the  student, 
as  it  might  result  in  an  explosion  in  unskilled  hands. 

NaClO4+H2SO4->NaHSO4+HC104. 

Perchloric  acid  is  a  colorless  liquid.  It  is  a  violent  oxidizing 
agent,  as  shown  by  the  fact  that  a  drop  of  the  acid  will  set  fire  to 
filter  paper.  The  diluted  acid  is  now  coming  into  use  in  labo- 
ratories as  an  oxidizing  agent,  and  also  for  the  purpose  of  pre- 
cipitating potassium  perchlorate  in  the  quantitative  analysis 
of  potassium. 


CHAPTER  XVI 

HEAT  AND  ENERGY 

356.  Heat  of  Combustion. — Since  coal,  wood,  and  fuel  gas 
are  burned  ordinarily  in  order  to  produce  heat  rather  than  as 
a  means  of  obtaining  their  products  of  combustion,  carbon 
dioxide  and  water,  it  becomes  a  matter  of  importance  to  discover 

how  much  heat  is  produced 
in  the  burning  of  a  known 
weight  of  a  given  substance. 
The  unit  of  heat  is  the 
calorie  (in),  which  is  the 
amount  of  heat  required  to 
raise  the  temperature  of  one 
gram  of  water  one  degree  centi- 
grade. The  amount  of  heat 
produced  by  the  burning  of 
one  formula  weight  of  a  pure 
substance  is  called  its  heat 
of  combustion.  The  heat  of 
combustion  of  a  solid  is  de- 
termined by  burning  a  known 
weight  of  it  within  an  appa- 
ratus of  special  design,  called 
a  bomb  calorimeter. 

357.  The  Bomb  Calorimeter. — This  apparatus,  illustrated  in 
Fig.  44,  consists  of  a  heavy-walled  metallic  bomb  with  a  gas- 
tight  cover,  surrounded  by  a  vessel  of  water.    The  latter  is  con- 
tained in  a  larger  vessel  with  walls  of  heat-insulating  material. 
A  weighed  amount  of  substance  whose  heat  of  combustion  is 
to  be  found  is  placed  in  the  crucible  of  the  bomb,  which  is  filled 
with  oxygen  gas.    The  substance  is  then  ignited  by  heat  from 
a  wire  which  carries  an  electric  current.    The  temperature  of 
the  water  surrounding  the  bomb  is  measured  accurately  before 


FIG.  44 


214 


Heat  and  Energy 


and  after  the  burning,  and  the  number  of  calories  of  heat  pro- 
duced is  calculated  from  the  rise  of  temperature  and  the  weight 
of  water  actually  heated,  plus  the  water  equivalent  of  the  bomb, 
etc.  The  water  equivalent  is  the  amount  of  water  which  has 
the  same  heat  capacity  as  the  bomb  and  other  heated  parts  of 
the  apparatus.  Some  typical  results  of  measurements  of  heats 
of  combustion  are  shown  in  Table  XI.  The  values  are  given 
to  the  nearest  hundred,  since  this  is  about  the  limit  of  accuracy 
in  such  measurements. 

TABLE  XI 


Substance 

Calories  per  gram 

Formula 

Heat  of  Combustion 

Carbon  
Hydrogen 

8,130 
•jj.  4.00 

C  =12  g. 
H2=     2g 

97,000 

68  800 

Sulfur 

2  2OO 

S  =12  S 

7O  4.OO 

Acetylene 

II,9OO 

C2Ha=26g 

•21  r  4.00 

Carbon  monoxide  

2,430 

C0=28g. 

68,200 

Since  one  formula  weight  of  a  gaseous  substance  has  a  volume 
22.4  liters,  the  heats  of  combustion  of  H2,  C2H2,  and  CO  are 
the  amounts  of  heat  produced  in  the  burning  of  equal  volumes 
of  these  gases.  It  will  be  seen  that  the  heat  of  combustion  of 
C2H2  is  very  large  (nearly  five  times  that  of  hydrogen).  This 
accounts  in  part  for  the  very  high  temperature  of  the  oxyacety- 
lene  flame  (315). 

358.  The   British   Thermal   Unit,    B.T.U. — In   engineering 
practice  quantities  of  heat  are  measured  in  British  Thermal 
Units  (B.T.U.)  instead  of  in  calories.     This  unit  is  the  amount 
of  heat  required  to  raise  the  temperature  of  one  pound  of  water  one 
degree  Fahrenheit.    Since  one  pound  equals  453  g.,  and  i°  F.  =  5/9 
of  i°  C.,  it  follows  that  i  B.T.U. =252  calories.    The  heat  pro- 
duced in  burning  coal,  coke,  and  fuel  gas  is  called  its  calorific 
power.     It  is  usually  stated  in  terms  of  B.T.U.  per  pound  of 
fuel. 

359.  Composition  and  Calorific  Power  of  Fuel. — Since  the 
value  of  fuel  is  directly  dependent  on  its  calorific  power,  the 
testing  of  fuel  is  a  matter  of  great  practical  importance.    In 
testing  coal  it  is  customary  to  determine  the  moisture,  volatile 


2l6 


Introduction  to  General  Chemistry 


matter,  "fixed  carbon,"  and  ash  in  addition  to  the  calorific 
power.  The  "fixed  carbon' '  is  the  non- volatile  residue  left 
when  all  volatile  matter  is  driven  off  at  a  bright-red  heat  in  the 
absence  of  air,  less  the  ash  contained  therein.  The  calorific 
power  is  usually  expressed  in  terms  of  B.T.U.  per  pound  of  fuel, 
or  per  cubic  foot  in  the  case  of  gases.  Table  XII  gives  some 
results  for  a  variety  of  solid  fuels. 

TABLE  XII 


KIND  OF  FUEL 

PERCENTAGE  COMPOSITION 

CALORIFIC  POWER 

Volatile 
Matter 

Fixed 
Carbon 

Ash 

Calories 
per  gram 

B.T.U. 
per  pound 

Lacka  wanna  anthracite  coal 
Pocahontas  coal  

18 
35 
o-5 

38 

84 
74 
50 
90 

Si 

ii 

7 
6 

9 
4 
0.4 
0.4 

7,724 
8,760 
8,080 
7,900 
7,200 
4,600 
5,000 
11,520 

13,900 
15,680 
14,540 
14,200 
13,000 
8,300 
9,100 
20,736 

Indiana  bituminous  coal  .  .  . 
Coke 

Lignite 

Oak  wood 

Pine  wood  (resinous)  .... 

Crude  petroleum 

Table  XIII  gives  the  calorific  power  of  some  typical  fuel 

gases. 

TABLE  XIII 

CALORIFIC  POWE^IN  B.T.U.  PER  CUBIC  FOOT 

Kokomo,  Indiana,  natural  gas i  ,000 

Pittsburgh,  Pennsylvania,  natural  gas 1,150 

Coal  gas 650 

City  of  Chicago  gas 600 

360.  The  Evaporation  of  Water  and  the  Production  of 
Steam. — We  can  easily  calculate  the  amount  of  fuel  theoretically 
needed  to  change  water  at  ordinary  temperature  into  steam. 
If  one  gram  of  water  at  20°  is  heated  to  100°,  80  calories  of  heat 
are  required,  and  in  addition  540  calories  are  needed  to  change 
this  hot  water  into  steam.  The  total  is  620  calories.  Since 
the  burning  of  one  gram  of  coal  produces  about  8,000  calories, 
if  all  this  heat  were  utilized  it  would  be  sufficient  to  evaporate 
(8,000-:-  620)  13  g.  of  water.  In  practice  much  heat  is  lost  to 
the  surroundings,  as  well  as  in  the  hot  smoke  which  goes  up  the 


Heat  and  Energy  217 

smokestack.  Engineers  consider  that  it  is  good  practice  to 
evaporate  8  g.  of  water  with  i  g.  of  coal.  Therefore  one  pound 
of  good  coal  will  change  8  lb.,  or  about  i  gal.  (8.3  Ib)  of  water 
at  ordinary  temperature  into  steam  at  100°. 

361.  Heat  of  Reaction  and  Heat  of  Formation. — We  have 
already  frequently  observed  that  numerous  reactions  other  than 
combustions  in  oxygen  (air)  produce  much  heat.    Among  such 
are  the  reactions  of  chlorine  with  hydrogen  (244),  phosphorus 
(247),  antimony  (246),  and  turpentine  (248);   and  water  with 
sulfuric  acid  (93),  potassium  (106),  and  calcium-  oxide  (150). 
The  heat  produced  in  these  and  other  reactions  may  be  measured 
in  suitably  constructed  calorimeters  and  the  results  expressed 
most  conveniently  by  stating  the  amount  of  heat  given  out  in 
the  reaction  of  formula  weights  of  the  uniting  substances;   or 
in  the  formation  of  one  formula  weight  of  the  product.     Thus 
the  heat  of  reaction  of  CaO  and  H2O  may  be  written 

CaO+H2O->Ca(OH)2+5,ioo  cal. 

and  the  heat  of  formation  of  water  from  its  elements 
H2+JO3->H2O+ 68,800  cal. 

362.  Heat  of  Neutralization. — The  union  of  acids  and  bases 
to  form  salts  and  water  always  gives  out  heat.     In  fairly  dilute 
solutions  the  amount  of  heat  given  out  when  one  formula  weight 
of  water  is  so  formed  is  almost  exactly  the  same  for  many  acids 
and  bases.     For  example. 

HC1,     NaOH  =  13,700  cal. 
HC1,     KOH  =13,700  " 
HNO3,  NaOH  =  13,700  " 
HNO3,  KOH  =13,700  " 

This  regularity  is  indeed  striking  and  must  mean  close  similarity 
in  the  processes  of  these  reactions.  How  chemists  interpret  this 
phenomenon  will  be  considered  in  chapter  xviii. 

363.  The  Law  of  Constant  Heat  Summation. — Let  us  now 
consider  the  following  question:   If  equal  quantities  of  a  given 
substance  can  be  changed  into  the  same  product  by  two  different 


218  Introduction  to  General  Chemistry 

ways,  will  the  amounts  of  heat  produced  be  the  same  in  the  two 
cases?  Carbon,  for  example,  gives  carbon  dioxide  when  it  is 

burned, 

C+O2->CO2, 

but  in  a  deficiency  of  oxygen  the  product  is  carbon  monoxide, 

2C+O2->2CO. 

Carbon  monoxide  is  a  colorless  gas  which  burns  readily,  giving 

carbon  dioxide, 

2CO+O2->2CO2. 

Therefore  it  is  possible  to  change  given  weights  of  carbon  and 
oxygen  into  carbon  dioxide  in  two  different  ways.  The  heats 
of  combustion  are  as  follows: 

FIRST  WAY 

C+|Oa->CO  +29,400  cal. 
CO+£O2-»CO2+68,2oo  cal. 


Sum    97,600  cal. 

SECOND  WAY 
C+O2->CO2+97,6oo  cal. 

These  results  show  that  if  1 2  g.  of  carbon  (C  =  1 2)  unite  with 
32  g.  of  oxygen  (02  =  32  and  J02  =  16)  the  total  heat  produced  is 
the  same  no  matter  in  which  way  the  union  occurs. 

Another  illustration  is  found  in  the  formation  of  a  solution 
of  ammonium  chloride,  NH4C1,  from  NH3  and  HC1  gases. 
This  reaction  can  take  place  in  two  ways: 

FIRST  WAY 

NH3(gas)+HCl  (gas)->NH4Cl  (solid) +42,000  cal. 
Heat  absorbed  in  dissolving  the  NH4C1  in  water  =  —  3,900  cal. 


Excess  of  heat  produced  over  heat  absorbed  =  38,100  cal. 

SECOND  WAY 

Heat  of  solution  of  NH3  in  water                 =  8,400  cal. 

Heat  of  solution  of  HC1  in  water                 =  17,300  cal. 

Heat  of  neutralization  of  the  two  solutions  =  12,400  cal. 


Total  heat  produced  =     38,100  cal. 


Heat  and  Energy  219 

Innumerable  cases  like  the  two  here  given  in  illustration  have 
led  to  the  Law  of  Constant  Heat  Summation  (Law  of  Hess). 
The  heat  produced  or  absorbed  in  the  change  of  given  substances 
into  the  same  final  products  (in  the  same  physical  state)  is  the  same, 
by  whatever  way  the  changes  occur. 

That  the  heat  of  a  given  reaction  is  dependent  on  the  physical 
state  of  the  reacting  substances  and  products  is  illustrated  by 
the  following  example: 

CaO-f  H2O  (liquid)->Ca(OH)2  (solid) +15,100  cal. 
CaO+H2O  (ice)      ->Ca(OH)2  (solid) +13, 700  cal. 

Difference  =1,400  cal. 

The  difference,  1,400  cal.,  is  due  to  the  fact  that  it  requires  this 
amount  of  heat  to  change  one  formula  weight  of  ice  into  water 
(18X79  =  1,422)  (118). 

364.  Heat  Produced  in  Slow  Oxidation.  Spontaneous  Com- 
bustion.— Numerous  experiments  have  proved  that  the  amount 
of  heat  formed  in  a  given  reaction  is  just  the  same  whether  the 
change  takes  place  slowly  or  rapidly.  The  decay  of  wood  leads 
ultimately  to  the  production  of  carbon  dioxide  and  water,  the 
same  products  as  those  formed  when  wood  is  burned.  During 
the  decay  of  wood,  heat  is  produced  so  slowly  that  its  formation 
is  usually  not  perceptible  by  ordinary  observation.  Coal  also, 
when  exposed  to  the  air,  slowly  oxidizes.  In  so  doing  it  often 
loses  an  appreciable  part  of  its  heating  value  before  it  is  burned. 
The  depreciation  on  this  account  in  the  value  of  stored  coal  is  a 
matter  of  considerable  importance. 

If  coal  (especially  bituminous  coal)  in  small  lumps  and  con- 
taining much  dust  is  heaped  in  immense  piles,  such  as  are  seen 
in  coal  yards,  the  heat  produced  by  the  slow  oxidation  does  not 
escape  readily  from  the  bottom  layers  of  the  pile.  The  result 
is  a  gradual  rise  of  temperature.  At  the  higher  temperature 
oxidation  and  therefore  heat  production  go  on  still  faster,  since 
usually  enough  air  can  diffuse  in  to  keep  up  the  supply  of  oxygen. 
Finally  the  temperature  may  rise  so  high  that  the  pile  of  coal 
actually  takes  fire  at  the  surface,  where  there  is  of  course  an 
unlimited  supply  of  oxygen.  Fire  originating  in  this  way  is 


22O  Introduction  to  General  Chemistry 

said  to  be  due  to  spontaneous  combustion.  The  loss  of  coal 
through  such  fires  was  a  very  serious  feature  of  the  "coal  famine" 
of  1917-18.  Some  smoke  is  seen  issuing  from  the  majority  of 
large  piles  of  low-grade  coal  in  the  Chicago  district,  thus  indi- 
cating more  or  less  fire  beneath.  It  is  almost  impossible  to 
extinguish  fire  in  a  very  large  coal  pile.  The  best  way  to  prevent 
serious  rise  of  temperature  in  coal  piles  is  to  provide  numerous 
air  shafts  in  the  pile,  by  means  of  which  warm  air  can  escape. 
This  does  not  entirely  prevent  oxidation  but  keeps  the  tempera- 
ture down  to  a  point  where  the  oxidation  is  not  dangerously 
fast. 

It  is  a  popularly  known  fact  that  "greasy"  rags  will  often 
catch  fire  spontaneously.  As  a  matter  of  fact  such  fires  originate 
usually  in  rags  soaked  in  oils  used  in  paint  or  varnish,  especially 
linseed  oil  or  turpentine.  The  "drying"  of  paint  and  varnish 
is  not  a  process  of  evaporation  as  much  as  one  of  oxidation  of 
the  oil  used.  These  paint  and  varnish  oils  readily  unite  with 
oxygen  to  form  solid  products.  In  this  process  heat  is  produced. 
In  a  pile  of  rags,  etc.,  covered  with  such  oils  sufficient  rise  of 
temperature  may  occur  to  cause  spontaneous  combustion.  For 
this  reason  greasy  rags,  etc.,  should  never  be  left  where  they 
can  do  damage  if  they  take  fire. 

365.  Dust  Explosions. — When  the  air  is  filled  with  the  dust  of 
coal,  wood,  flour,  or  other  combustible  substance  a  flame  will 
often  start  a  combustion  which  will  spread  with  explosive 
rapidity.    Appalling  explosions  have  occurred  from  such  causes 
in  coal  mines,  wood-working  factories,  and  flour  mills.     Even 
dust  which  is  at  rest  in  such  places  is  blown  into  the  air  by  the 
on-coming  explosion  wave  and  is  thus  changed  to  an  explosive 
dust  and  air  mixture.     It  is  easy  to  see  that  a  dust  explosion  is 
due  to  the  extremely  rapid  burning  of  minute  particles,  each 
surrounded  by  an  abundance  of  oxygen.     Dust  explosions  are 
best  prevented  by  keeping  mines,  mills,  etc.,  free  from  accumula- 
tions of  dust. 

366.  Modes  of  Heat  Production  in  Physical  and  Chemical 
Changes. — We  have  now  learned  that  heat  is  produced  (or 
absorbed)  in  a  variety  of  physical  and  in  all  chemical  changes. 


Heat  and  Energy  221 

The  following  seven  modes  of  heat  production  (or  absorption) 
have  been  studied: 

1.  Latent  heat  of  fusion  (melting)  (118). 

2.  Latent  heat  of  evaporation  (115). 

3.  Heat  of  solution  (127). 

4.  Heat  of  combustion  (356). 

5.  Heat  of  formation  (361). 

6.  Heat  of  reaction  (361). 

7.  Heat  of  neutralization  (362). 

The  first  three  modes  have  to  do  with  physical  changes  of  the 
sort  known  as  changes  of  state ;  the  last  four  are  due  to  chemical 
changes.  All  changes  of  state  and  many  chemical  changes  are 
reversible  processes.  In  every  reversible  process,  if  heat  is 
given  out  when  the  change  proceeds  in  one  direction,  heat  is 
absorbed  in  equal  amount  when  the  change  proceeds  to  an  equal 
extent  in  the  opposite  direction.  A  change  which  results  in  the 
production  of  heat  is  called  an  exothermic  change;  one  in  which 
heat  is  absorbed  is  an  endothermic  change. 

367.  Heat  Production  and  Equilibrium. — In  chapter  xiii 
(288)  the  effect  of  temperature  on  equilibrium  was  discussed 
briefly.  With  respect  to  the  change  of  solubility  it  was  stated 
that  raising  the  temperature  causes  that  change  of  solubility  to 
occur  which  invokes  an  absorption  of  heat.  We  also  saw  (288) 
that  for  chemical  equilibrium  raising  the  temperature  causes 
the  equilibrium  to  shift  in  the  direction  that  involves  an  absorption 
of  heat.  These  laws  are  entirely  general  and  apply  to  all 
changes  of  state  and  all  chemical  changes. 

In  the  shift  of  equilibrium  which  occurs  with  change  of 
temperature  the  fraction  of  the  reacting  substances  transformed 
to  new  products  is  determined,  in  a  given  case,  by  the  change 
of  temperature  (measured  in  degrees).  The  amount  of  heat 
(in  calories)  absorbed  (if  the  temperature  is  raised)  or  given  out 
(if  the  temperature  is  lowered)  is  determined  by  the  amount  of 
material  transformed.  An  example  will  make  the  matter  clearer. 

Hydrogen  and  iodine  vapor  react  partially  in  the  neighbor- 
hood of  400°  to  give  hydrogen  iodide  (264,  281,  288): 
H2-fI2±52HI+i,ooocal. 


222  Introduction  to  General  Chemistry 

This  equation  means  that  the  formation  of  two  formula  weights 
of  HI  from  H2  and  I2  (vapor)  at  about  40x3°  takes  place  with 
the  liberation  of  1,000  cal.  of  heat  or  500  cal.  for  each  formula 
weight  of  HI  produced.  The  following  table  shows  the  propor- 
tions of  molecules  in  the  equilibrium  mixture  at  370°  and  440°: 

H,  I,  HI  Total 

370°        100        100        800        1,000 
440°        no        no        780        1,000 

We  see  that  if  the  temperature  is  raised  from  370°  to  440°,  20 
molecules  of  HI  out  of  a  total  of  1,000  molecules  (2  per  cent  of 
the  whole)  change  into  H2  and  I2.  If  the  total  amount  of 
material  in  the  mixture  is  that  resulting  from  one  formula  weight 
each  of  H2  and  I2  (equivalent  to  two  formula  weights  of  HI), 
and  if  2  per  cent  of  the  whole  number  of  molecules  change  into 
H2  and  I2,  the  heat  absorbed  is  0.02 X  1,000  cal.  =  20  cal. 

368.  Work  and  Energy. — The  terms  work  and  energy  have 
very  definite  meanings  in  science.  The  subject  of  physics  is 
largely  concerned  with  these  very  important  matters;  and  since 
it  is  assumed  that  the  student  has  already  studied  physics,  an 
elementary  discussion  of  these  very  important  topics  is  unneces- 
sary. We  may,  however,  briefly  summarize  some  of  the  more 
prominent  points.  The  typical  example  of  work  in  the  physical 
sense  is  the  lifting  of  a  weight.  The  scientific  unit  of  work  is 
the  gram  centimeter,  which  is  the  work  required  to  lift  one  gram 
one  centimeter.  The  amount  of  work  done  in  lifting  a  weight  is 
the  product  of  the  force  required  (which  in  this  case  is  equal 
to  the  weight  in  grams)  and  the  vertical  distance  measured  in 
centimeters.  Thus  the  lifting  of  600  g.  to  a  height  of  30  cm. 
requires  the  doing  of  600X30  =  1 8,000  g.cm.  of  work.  The 
weight  of  600  g.,  having  been  lifted  30  cm.,  could  do  work  to 
the  extent  of  18,000  g.cm.  in  descending  30  cm.  It  is  said  to 
have  the  power  to  do  this  amount  of  work.  Now  power  to  do 
work  is  called  energy,  and  therefore  it  has  18,000  g.cm.  of 
energy.  Two  kinds  of  energy  are  recognized:  potential  energy, 
as  possessed  by  a  weight  which  may  do  work  on  descending,  and 
kinetic  energy,  or  the  energy  of  a  body  in  motion.  It  requires 


Heat  and  Energy 


223 


work  to  set  a  body  in  motion,  and  conversely  a  body  in  motion 
is  able  to  do  work. 

369.  The  Mechanical  Equivalent  of  Heat.— Heat  is  also  a 
form  of  energy,  because  heat  is  able  to  do  work.  A  steam  engine 
is  merely  a  machine  which  converts  the  heat  of  burning  coal  into 
kinetic  energy.  The  change  of  kinetic  energy  into  heat  may  be 
observed  on  every  hand:  anything  that  restrains  or  stops  the 
motion  of  a  moving  body  converts  part  or  all  of  its  kinetic  energy 
into  heat.  We  measure  energy  in  gram  centimeters  and  heat 
in  calories,  and  if  heat  is  a  form  of  energy  then  the  calorie,  like 
the  gram  centimeter,  must  be  an  energy  unit.  It  will  at  once 


FIG.  45 

be  asked:  Do  these  units  represent  equal  amounts  of  energy? 
In  other  words,  will  one  gram  centimeter  of  work  produce  one 
calorie  of  heat?  If  not,  how  many  gram  centimeters  are  required 
to  produce  one  calorie?  This  question  was  first  answered  by 
Joule  in  1840. 

370.  Joule's  Experiment. — In  Joule's  experiment,  with 
apparatus  shown  in  Fig.  45,  a  weight,  W,  attached  to  a  cord 
wound  on  a  cylinder,  in  slowly  descending  turns  a  stirrer  which 
is  surrounded  by  water  in  a  calorimeter,  C.  The  water,  which 
restrains  the  motion  of  the  stirrer,  becomes  warmer,  owing  to 
the  change  of  work  into  heat.  The  amount  of  work  in  gram 
centimeters  done  in  heating  the  water  is  the  product  of  the  mass 
in  grams  of  the  weight  and  the  distance  of  its  descent  in 


224  Introduction  to  General  Chemistry 

centimeters.  The  amount  in  calories  of  heat  produced  is  the 
product  of  the  rise  in  temperature  in  degrees  C.  and  the  mass, 
in  grams,  of  water  plus  the  water  equivalent  of  the  heated  parts 
of  the  calorimeter. 

By  means  of  this  apparatus  Joule  found  pretty  closely  the 
number  of  gram  centimeters  of  work  equivalent  to  one  calorie 
of  heat.  More  refined  work  since  then  has  shown  that  one 
calorie  is  equal  to  42,700  g.cm.  This  ratio  is  called  the  mechanical 
equivalent  of  heat.  This  means,  for  example,  that  one  gram 
falling  42,700  cm.  (a  little  over  a  quarter  of  a  mile)  produces  one 
calorie. 

371.  The  Conservation  of  Energy. — At  the  time  Joule  began 
his  experiments  in  1840  it  was  not  at  all  clear  that  the  amount 
of  heat  produced  by  a  given  amount  of  work  (kinetic  or  potential 
energy)  was  definite.  It  seemed  possible,  if  not  probable,  that 
different  modes  of  changing  work  into  heat  would  give  different 
values  for  the  mechanical  equivalent.  So  Joule  used  not  only 
the  method  and  apparatus  already  described  but  also  two  others. 
His  three  methods  and  the  mechanical  equivalent  of  one  calorie 
were  as  follows:  (i)  stirring  water  in  a  brass  vessel  with  a  brass 
paddle,  42,400  g.cm.;  (2)  stirring  mercury  in  an  iron  vessel 
with  an  iron  paddle,  42, 500 g.cm.;  (3)  rubbing  two  iron  rings 
together  under  mercury,  42,500  g.cm. 

The  very  close  agreement  of  the  results  of  the  three  experi- 
ments led  Joule  to  conclude  that  the  amount  of  heat  produced 
by  a  given  amount  of  work  is  always  the  same,  by  whatever  way  the 
work  is  changed  into  heat.  This  result  has  been  amply  confirmed 
by  all  later  experiments  and  experience.  When  work  of  any 
kind  (mechanical  energy,  either  kinetic  or  potential)  is  changed 
into  heat  there  is  no  real  loss  or  destruction  of  energy,  since  the 
heat  produced  is  also  energy  in  another  form  and  exactly  equal 
in  amount  to  the  work  done  in  producing  it.  This  conclusion 
is  concisely  stated  in  the  Law  of  the  Conservation  of  Energy: 
Energy  is  indestructible. 

Just  as  the  law  of  the  conservation  (indestructibility)  of 
matter  (21)  is  the  foundation  stone  of  the  science  of  chemistry, 
so,  similarly,  this  law  of  the  conservation  (indestructibility) 


Heat  and  Energy  225 

of  energy  is  the  solid  rock  upon  which  the  whole  structure  of  the 
science  of  physics  rests. 

372.  Other  Forms  of  Energy. — We  have  defined  the  term 
energy  as  the  power  of  doing  work;  and  since  heat  is  also  a  form 
of  energy,  we  might  extend  the  definition  so  as  to  read:   Energy 
is  the  power  to  do  work  or  produce  heat.    According  to  this  defini- 
tion of  energy  it  is  obvious  that  light  and  even  sound  and  espe- 
cially electric  currents  are  also  forms  of  energy,  since  each  of 
these  by  appropriate  means  can  produce  work  or  heat. 

373.  Chemical  Energy. — For  the  chemist  an  important  ques- 
tion now  arises:  What  shall  be  said  of  the  source  of  energy  that 
produces  the  great  heat  of  a  burning  substance?     This  question 
is  somewhat  like  the  one,  What  is  the  source  of  energy  of  a 
"  wound-up"  watch  spring?    To  wind  up  the  spring  a  certain 
amount  of  work  must  be  done.     Is  it  not  reasonable  to  say  that 
the  energy  used  in  winding  up  the  spring  has  been  "  stored  up" 
in  the  coiled  spring?     If  so,  we  may  say  that  this  energy  is 
changed  into  potential  energy,  just  as  we  say  that  the  energy 
required  to  lift  a  weight  is  changed  into  potential  energy  and 
can  be  regained  as  useful  work  then  the  weight  is  allowed  to 
descend.     Reasoning  somewhat  similarly,  we  may  conclude  that 
the  energy  given  out  as  heat  in  the  burning  of  hydrogen,  for 

which  we  have 

H2+JO2->H2O+68,8oo  cal., 

comes  from  some  form  of  potential  energy  which  has  been  stored 
up  in  the  two  gases.  This  conclusion  is  rendered  highly  probable 
by  reason  of  the  fact  that  by  means  of  an  electric  current  (elec- 
trical energy)  we  can  decompose  water  into  hydrogen  and  oxygen. 
Since  the  electrical  energy  disappears  and  very  little  heat  is 
formed,  we  may  very  reasonably  conclude  that  it  has  been 
changed  into  some  sort  of  potential  energy  stored  up  in  the  two 
gases  formed  from  the  water.  The  form  of  potential  energy 
stored  up  in  chemical  substances  and  liberated  when  they  react 
is  called  chemical  energy. 

374.  The  Sun  as  a  Source  of  Energy.— It  will  be  interesting 
to  trace  some  familiar  form  of  energy  through  various  trans- 
formations back  to  its  source.    Take,  as  an  example,  the  energy 


226  Introduction  to  General  Chemistry 

given  out  as  light  and  heat  by  an  electric  lamp.  The  energy 
comes  to  the  lamp  as  an  electric  current  having  electrical  energy. 
This  electrical  energy  was  produced  hi  a  dynamo  or  generator, 
the  armature  (the  moving  part)  of  which  was  turned  by  a  steam 
engine.  The  kinetic  energy  of  the  engine  was  derived  from  hot, 
compressed  steam  produced  from  water  by  the  burning  of  coal 
which  has  resulted  from  the  slow  transformation  of  vegetable 
matter.  , 

Plants  derive  nearly  all  of  their  substance  from  water  and 
the  carbon  dioxide  of  the  air  under  the  influence  of  the  light  and 
heat  of  the  sun.  A  great  deal  of  energy  is  taken  up  by  plants 
as  light  and  heat  and  is  stored  as  chemical  energy  in  the  sub- 
stances composing  them,  as  well  as  in  the  oxygen  which  is  set 
free  by  the  growing  plant.  Recapitulating,  we  see  that  the 
light  and  heat  from  the  sun  are  changed  by  growing  plants  into 
chemical  energy;  this  energy  is  largely  conserved  when  plants 
are  changed  into  coal.  When  the  coal  burns,  its  chemical 
energy,  supplemented  by  that  of  the  oxygen  of  the  air,  is  changed 
into  heat,  which  is  in  turn  changed  into  kinetic  energy  in  the 
steam  engine.  The  kinetic  energy  of  the  engine  is  then  changed 
by  a  dynamo  into  electrical  energy,  and  the  latter  produces  in 
the  lamp  heat  and  light. 


CHAPTER  XVII 

THE  IONIC  HYPOTHESIS 

375.  The  Ionic  Hypothesis. — This  chapter  will  treat  of  the 
properties  and  behavior  of  acids,  bases,  and  salts  and  aims  to 
show  how  a  supposition  called  the  ionic  hypothesis  furnishes  a 
satisfactory  explanation  of  many  facts. 

376.  The  Two  Parts  of  a  Salt. — It  must  have  been  noticed 
that  a  salt  is  made  up  of  two  parts,  the  metallic  or  basic  part  and 
the  non-metallic  or  acidic  part.    The  latter  may  be  an  element 
like  chlorine  in  sodium  chloride;   or  it  may  be  a  radical  (147) 
like  SO4,  which  is  contained  in  every  sulfate.    The  name  of  a 
salt  always  indicates  the  parts  of  which  it  may  be  considered  as 
being  made  up.    Thus  potassium  nitrate,  KNO3,  is  composed 
of  potassium  and  nitrate  radical,  NO3;  and  calcium  carbonate, 
CaC03,  of  calcium  and  carbonate  radical,  CO3. 

377.  The  Two  Parts  of  an  Acid. — Every  acid  may  also  be 
considered  as  made  up  of  two  parts,  one  of  which  is  hydrogen 
and  the  other  the  characteristic  acid  radical  of  that  acid.    For 
example,  sulfuric  acid,  H2SO4,  may  be  considered  to  consist  of 
hydrogen  and  sulfate  radical,  SO4;  and  phosphoric  acid,  H3P04, 
to  consist  of  hydrogen  and  phosphate  radical,  PO4.    For  this 
reason  S04  and  P04  may  be  called  acidic  radicals.    Dilute 
solutions  of  pronounced  acids  have  a  sour  taste.    Since  hydro- 
gen is  the  only  constituent  which  all  acids  have  in  common,  we 
may  reasonably  conclude  that  the  sour  taste  is  due  to  the  H 
radical. 

378.  The  Two  Parts  of  a  Base. — A  base  is  usually  the 
hydroxide  of  a  metallic  element,  'and  it  may  therefore  be  con- 
sidered as  made  up  of  two  parts,  the  metal  and  the  hydroxyl 
radical,    OH.    Thus   sodium   hydroxide,    NaOH,    consists   of 
sodium   and   hydroxyl   radical,   OH.    Ammonium   hydroxide, 
NH4OH,  may  be  considered  as  made  up  of  ammonium  radical, 
NH4,  and  hydroxyl.    Consequently  the  ammonium  radical  may 

227 


228  Introduction  to  General  Chemistry 

be  called  a  basic  radical.  It  is  the  only  basic  radical  that  we 
have  studied,  although  many  others  are  known. 

379.  The  Process  of   Neutralization. — The   two   following 
equations  represent  typical  cases  of  neutralization: 

NaOH+HCl->NaCl+H2O; 
NH4OH+HNO3-»NH4NO3+H2O. 

We  notice  that  in  each  case  the  salt  which  is  formed  is  made  up 
of  two  parts,  one  of  which  comes  from  the  base,  the  other  from 
the  acid.  The  remaining  parts  of  acid  and  base,  H  and  OH, 
unite  to  form  water.  We  might  call  water  hydrogen  hydroxide 
and  think  of  it  as  being  made  up  of  two  parts  hydrogen  and 
hydroxyl  radical.  The  process  of  neutralization  consists,  there- 
fore, merely  of  an  exchange  of  partners,  so  to  speak,  on  the  part 
of  the  base  and  the  acid. 

As  a  matter  of  fact,  not  only  can  neutralization  be  represented 
in  this  way,  but  most  reactions  in  water  solution  between  acids, 
bases,  and  salts  which  do  not  involve  oxidation  or  reduction 
may  be  regarded  as  an  exchange  of  the  partners  of  the  reagents 
initially  used.  This  will  be  made  clear  by  the  following  examples. 

380.  Reactions  of  Barium  Salts  with  Sulfates. — If  we  add 
dilute  sulfuric  acid  to  a  dilute  solution  of  barium  chloride  a  white 
precipitate  of  barium  sulfate  is  formed, 

H2S04+BaCl2->BaS04+2HCl. 

A  precipitate  of  barium  sulfate  also  results  when  a  solution 
of  any  barium  salt  is  added  to  a  solution  of  any  sulfate,  as 
illustrated  by  the  following  equations : 

K2SO4+Ba(NO3)2->BaSO4+2KNO3, 
CuSO4+BaBr2-»BaSO4-f-CuBr2. 

This  is  so  generally  true  that  the  formation  of  a  precipitate 
of  barium  sulfate  upon  the  addition  of  a  solution  of  a  barium 
salt  to  some  other  solution  shows  that  this  second  solution  con- 
tains the  sulfate  radical,  S04,  in  the  form  either  of  a  sulfate  or 
of  sulfuric  acid.  We  say  therefore  that  the  formation  of  a 
precipitate  of  barium  sulfate  when  a  solution  of  a  barium  salt  is 
added  to  another  solution  is  a  test  for  the  sulfate  radical.  It  is 


The  Ionic  Hypothesis  229 

important  to  note  that  it  is  the  SO4  radical,  and  not  sulfur  or 
oxygen  or  a  combination  of  the  two  in  some  other  proportion, 
that  responds  to  this  test.  A  solution  of  hydrogen  sulfi.de, 
H2S  (339),  which  may  be  considered  as  being  made  up  of  two 
parts,  hydrogen  and  sulfur,  does  not  give  a  precipitate  of  any 
sort  with  a  solution  of  a  barium  salt.  Furthermore,  pure  dilute 
sulfurous  acid,  H2S03  (340),  which  is  made  up  of  hydrogen  and 
sulfite  radical,  SO3,  does  not  give  a  precipitate  when  mixed  with 
a  barium  salt  solution. 

381.  Reactions  of  Simple  Lead  Salts.— Lead  sulfate,  PbS04, 
is  also  a  white  insoluble  salt.    If  we  add  a  solution  of  any 
simple  lead  salt  to  a  dilute  solution  of  sulfuric  acid  or  any  soluble 
sulfate,  we  obtain  a  white  precipitate  of  lead  sulfate, 

Na2SO4+Pb(N03)2^PbS04+2NaNO3. 

In  this  case,  just  as  in  the  precipitation  of  barium  sulfate,  it  is 
the  sulfate  radical,  S04,  which  has  united  with  the  lead  to  form 
the  precipitate. 

It  is  also  of  equal  interest  to  note  that  if  the  nitrate  of  barium 
or  of  lead  is  used,  the  nitrate  radical,  NO3,  is  left  in  combination 
with  the  basic  element  or  radical  which  was  originally  combined 
with  the  sulfate  radical. 

382.  The  Reaction  of  Silver  Salts  with  Chlorides.— We  have 
already  learned  that  a  solution  of  silver  nitrate  reacts  with 
hydrochloric  acid  or  a  chloride  to  give  a  white  precipitate  of 
silver  chloride : 

AgNO3+NaCl^AgCl+NaNO3.  (169) 

A  solution  of  any  simple  silver  salt  reacts  similarly  with  hydro- 
chloric acid  or  any  chloride,  so  that  we  may  think  of  the  reaction 
as  characteristic  and  call  it  a  test  for  silver  salts.  This  reaction 
is  specifically  that  of  the  chloride  radical;  for  if  we  add  silver 
nitrate  solution  to  a  solution  of  potassium  chlorate,  KC103 
(353),  no  apparent  change  is  observed;  certainly  no  silver 
chloride  is  formed,  otherwise  the  latter,  being  insoluble,  would 
separate  out  as  a  white  precipitate.  This  shows  that  chlor- 
ine in  the  chlorate  radical,  C1O3,  behaves  entirely  differently 


230  Introduction  to  General  Chemistry 

from  chlorine  in  the  form  of  a  chloride.  We  also  find  that 
solutions  of  perchlorates,  of  which  potassium  perchlorate, 
KC1O4  (355),  is  an  example,  do  not  give  precipitates  with  solu- 
tions of  silver  salts.  It  is  possible  to  make  both  silver  chlorate, 
AgClO3,  and  silver  perchlorate,  AgClO4,  by  methods  which  we 
need  not  consider  at  present,  and  it  is  found  that  these  salts  are 
entirely  different  from  silver  chloride,  and  that  both  are  easily 
soluble  in  water. 

This  brief  review  of  reactions,  most  of  which  have  already 
been  studied,  is  sufficient  to  illustrate  the  subject  in  hand,  but 
many  other  examples  of  the  same  principle  will  be  found  in  the 
previous  chapters. 

383.  Summary  and  Conclusions. — The  observations  which 
we  have  made  are  typical  for  all  acids,  bases,  and  salts.    Each 
may  be  shown  to  be  made  up  of  two  parts.    In  the  examples  we 
have  studied  these  are,  on  the  one  hand,  hydrogen,  a  metal,  or 
the  ammonium  radical,  and,  on  the  other,  hydroxyl,  a  halogen, 
sulfur,  or  an  acid  radical.    Hydrogen  is  one  of  the  two  parts  of 
every  acid,  and  hydroxyl  one  of  the  two  parts  of  every  base. 
In  chemical  reactions  between  acids  and  bases,  acids  and  salts, 
bases  and  salts,  and  between  two  salts  (where  oxidation  and 
reduction  do  not  occur)  the  two  substances  simply  exchange 
parts.    This  kind  of  chemical  change  is  called  double  decompo- 
sition (337)  or  metathesis.     The  chemical  reactions  which  acids, 
bases,  and  salts  give  are  in  reality  only  the  reactions  of  their  parts. 

Finally  it  should  be  noted  that  the  recombination  of  these 
parts  always  takes  place  between  the  basic  or  metallic  part  on 
the  one  hand  and  the  acidic  or  the  non-metallic  part  on  the 
other.  Double  decompositions  in  water  solutions  never  give 
compounds  such  as  KNa  or  C1S04.  This  is  a  striking  observa- 
tion, and  the  fact  that  we  have  as  yet  no  explanation  for  it 
warns  us  at  once  that  we  must  go  farther  in  our  observations 
to  understand  even  the  most  commonplace  of  these  reactions. 

384.  Double  Decomposition  and  Electrical  Conductivity. — 
Along  with  the  ability  to  undergo  double  decompositions,  acids, 
bases,  and  salts  in  their  water  solutions  have  the  property  of 
conducting  the  electric  current.    If  we  set  up  a  battery,  a 


The  Ionic  Hypothesis 


231 


galvanometer,  and  a  salt  solution  in  the  manner  shown  in 
Fig.  46,  using  platinum  electrodes  and  a  sufficient  number  of 
dry  cells  or  other  source  of  current  to  give  a  suitable  deflection 
of  the  galvanometer,  we  shall  find  that  if  we  replace  the  solution 
by  distilled  water  practically  no  current  will  be  indicated  by  the 
galvanometer.  We  also  find  that  if  dry  salt  is  substituted  for 
the  solution  no  current  will  pass.  If  now  we  pour  distilled 
water  on  the  salt  while  the  latter  is  still  in  contact  with  the 
electrodes,  a  current  begins  to  pass  through  the  solution  of  salt 
which  is  quickly  formed. 


J 


FIG.  46 

These  results  lead  to  the  conclusion  that  neither  pure  water 
nor  dry  salt  conducts  the  current  appreciably  compared  with 
the  solution  formed  from  salt  and  water.  All  other  soluble 
salts  behave  similarly.  It  is  also  easily  discovered  by  experi- 
ment that  dry  (water-free)  bases  and  acids  are  no  better  con- 
ductors  than  dry  salts,  although  water  solutions  of  acids  and 
bases  are  good  conductors.  Water  solutions  of  other  substances 
than  acids,  bases,  and  salts  do  not  conduct  electricity  any  better 
than  does  pure  water. 

The  close  connection  which  always  exists  between  electrical 
conductivity  and  the  ability  of  a  mixture  to  undergo  double 
decomposition  is  illustrated  by  the  following  experiment.  Ferric 
sulfate  and  sodium  carbonate  can  be  mixed  dry  without  any 
apparent  change;  but  let  the  mixture  once  be  wet  with  water, 
immediately  a  violent  evolution  of  gas  follows  and  a  red 
precipitate  of  ferric  hydroxide  appears.  That  the  mixture  of 
the  dry  substances  is  a  non-conductor  is  shown  by  placing  it 
in  the  dry  beaker,  Fig.  46.  No  current  passes,  but  when  water 
is  added  the  substances  dissolve,  and  the  solution  so  formed 
conducts  the  current.  At  the  same  time  the  chemical  reaction 


232  Introduction  to  General  Chemistry 

begins  vigorously.  Since  chemical  reactivity  and  electrical  con- 
ductivity seem  therefore  to  go  hand  in  hand,  we  shall  next  study 
the  behavior  of  solutions  of  acids,  bases,  and  salts  when  an  electric 
current  is  passed  through  them. 

385.  The  Electrolysis  of  Solutions. — We  have  already  learned 
that  the  electrolysis  of  concentrated  hydrochloric  acid  sets  free 
hydrogen  and  chlorine  (43),  and  that  the  electrolysis  of  common 
salt  (238)  yields  these  same  gases  and  in  addition  forms  sodium 
hydroxide.  In  the  case  of  hydrochloric  acid,  electrolysis  simply 
causes  the  separation  of  the  two  constituents, 

HC1->H+C1. 

On  being  set  free  the  atoms  of  the  two  elements  each  form 
diatomic  molecules,  thus, 

2H->H2,  and  2C1->C12. 

These  last  reactions  are  doubtless  secondary;  and  for  the  sake 
of  brevity,  in  the  examples  of  electrolysis  that  follow,  reactions 
of  this  kind  will  be  indicated  by  separate  equations  without 
further  comment. 

In  the  case  of  the  electrolysis  of  common  salt  it  seems  pos- 
sible, as  already  explained  (238),  that  the  first  change  is  a 
separation  into  sodium  and  chlorine,  thus : 

NaCl->Na+Cl, 

2C1->C12. 

The  sodium  then  reacts  with  the  water  present  to  form  sodium 
hydroxide  and  hydrogen: 

Na+H2O->NaOH-hH, 
2H->H2. 

Whether  this  is  the  only  possible  explanation  of  the  way  the 
changes  take  place  can  best  be  discussed  later;  but  it  can  be 
pointed  out  here  that  the  foregoing  equation  would  demand  that 
the  sodium  hydroxide  should  be  formed  at  the  electrode  at  which 
the  hydrogen  is  given  off;  and  this  is,  in  fact,  the  case.  When 
hydrogen  is  set  free  in  the  electrolysis  of  any  substance  it  always 


The  Ionic  Hypothesis  233 

appears  at  the  negative  electrode  or  cathode,  while  chlorine  is 
liberated  only  at  the  positive  electrode  or  anode. 

386.  The  Electrolysis  of  Copper  Salts.—  If  a  solution  of 
cupric  chloride,  CuCl2,  is  electrolyzed  between  platinum  poles 
or  electrodes,  copper  is  deposited  on  the  negative  pole  and 
chlorine  gas  is  set  free  at  the  positive  pole.  Here  again,  as  in 
the  case  of  hydrochloric  acid,  we  have  the  simplest  possible 
kind  of  a  change,  as  represented  by  the  following  equation: 

CuCla-»Cu+2Cl, 

2C1->C12. 

If  copper  sulfate,  CuS04,  is  electrolyzed,  a  plating  of  metallic 
copper  is  again  formed  on  the  negative  electrode,  while  from  the 
positive  electrode  oxygen  gas  is  given  off.  Examination  of 
the  products  after  electrolysis  shows  that  sulfuric  acid  is  con- 
tained in  the  solution  surrounding  the  positive  electrode.  In 
fact,  if  the  electrolysis  is  continued  until  all  the  copper  in  the 
original  solution  is  deposited,  all  the  sulfate  radical  of  the 
original  copper  sulfate  is  changed  into  sulfuric  acid,  and  this 
acid  is  contained  in  the  part  of  the  solution  surrounding  the 
positive  electrode.  The  formation  of  sulfuric  acid  and  oxygen 
may  be  explained  by  supposing  the  copper  sulfate  to  be  separated 
by  the  electric  current  into  copper  and  sulfate  radical,  S04,  and 
that  the  latter  reacts  with  water  to  form  sulfuric  acid  and  oxygen: 

SO4+H2O->H2S04+O, 


387.  The  Electrolysis  of  Silver  Nitrate.  —  If  an  electric  cur- 
rent is  passed  through  a  solution  of  silver  nitrate,  AgNO3,  silver  is 
deposited  on  the  negative  electrode  and  oxygen  and  nitric  acid 
appear  at  the  positive  electrode.    Probably  silver  nitrate  is 
first  separated  into  silver  and  nitrate  radical,  NO3;   the  latter 
then  reacts  with  water  to  form  nitric  acid  and  oxygen: 

2N03+H20->2HNO3+0, 

20->O2. 

388.  Summary.  —  In  Table  XIV  we  have  summarized  the 
results  just  cliscussed,  leaving  out  of  consideration  the  secondary 


234  Introduction  to  General  Chemistry 

changes  that  often  take  place  between  the  substance  set  free  and 
the  water.  We  see  that  the  parts  into  which  a  substance  is  separated 
by  electrolysis  are  the  same  as  those  which  change  partners  in  double 
decomposition  reactions. 

TABLE  XIV 
IMMEDIATE  PRODUCTS  OF  ELECTROLYSIS 


Liberated  at 
Negative  Plate 

Original 
Substance 

Liberated  at 
Positive  Plate 

H 

HC1 

Cl 

Na 

NaCl 

Cl 

Cu 

CuCl2                   *2C1 

Cu 

CuS04 

S04 

Ag 

AgNOj                    N03 

389.  The  Terms  Used  in  Electrolysis. — The  phenomena  of 
electrolysis  were  very  carefully  studied  about  1834  by  Faraday, 
who,  as  we  shall  see,  discovered  many  important  facts.    It  was 
Faraday  also  who  invented  the  terms  electrolysis,  electrolyze, 
electrode,  anode,  and  cathode.    He  called  the  solution  the 
electrolyte,  but  at  present  we  use  this  term  to  mean  the  dissolved 
substance.    That  part  of  the  electrolyte  which  during  electrolysis 
is  deposited  or  set  free  at  the  anode  or  positive  electrode  he 
called  the  anion.     The  other  part,  which  goes  to  the  cathode,  he 
called  the  cathion.    Frequently  he  had  occasion  to  speak  of 
anions  and  cathions  together,  and  then  he  referred  to  them  as 
the  ions  of  the  electrolyte.    For  example,  the  ions  of  copper 
sulfate  are  said  to  be  copper  and  sulfate  radical,  because  in 
electrolysis  copper  passes  to  and  is  deposited  on  one  electrode, 
while  the  sulfate  radical  goes  to  the  other.    Of  course  at  some 
time  or  other  the  radicals  or  the  partners  of  the  original  elec- 
trolyte must  have  broken  apart,  otherwise  they  could  not  have 
arrived  at  poles  distant  from  each  other. 

390.  Hydrogen  and  Metals  as  Cathions. — We  may  next  con- 
sider how  the  composition  of  the  ions  of  a  substance  can  be 
discovered. 

In  the  case  of  such  a  simple  substance  as  HC1  it  is  obvious 
that  the  ions  are  hydrogen  and  chlorine,  hydrogen  being  the 
cathion  and  chlorine  the  anion.  Since  all  acids  upon  electrolysis 


The  Ionic  Hypothesis  235 

give  off  hydrogen  at  the  cathode,  we  may  conclude  that  hydrogen 
is  the  cathion  of  all  acids. 

When  salts  are  electrolyzed  the  metal  is  either  deposited  in 
metallic  form  on  the  cathode,  as  in  the  case  of  copper  and  silver 
salts,  or  it  collects  in  the  solution  about  the  cathode  in  the  form 
of  hydroxide,  as  when  common  salt  is  used.  These  facts  lead 
to  the  conclusion  that  the  basic  or  metallic  elements  of  salts  are 
cathions. 

391.  Acid  Radicals  as  Anions. — On  the  other  hand,  the  acid 
elements  of  radicals  of  acids  and  salts  accumulate  at  the  anode 
and  are  either  given  off  as  free  elements,  as  in  the  case  of  chlorine, 
bromine,  and  iodine,  or  they  react  with  water  to  form  acids  and 
oxygen,  as  in  the  case  of  sulfate  and  nitrate  radicals. 

392.  Ions  and  Chemical  Reactions. — It  would  thus  appear, 
from  what  has  just  been  stated,  that  the  ions  of  an  acid  or  salt 
are  the  same  as  the  two  parts  of  which  its  chemical  reactions 
show  it  to  be  composed.     It  may  be  added  that  there  is  good 
reason  for  thinking  that  the  same  statement  also  applies  to 
bases.    The  cathion  of  a  base  is  usually  a  metal;    the  anion 
is  the  hydroxyl  radical. 

393.  Positive  and  Negative  Ions. — The  cathode  is  the  electro- 
negative electrode ;  to  it  go  the  cathions.     Since  it  is  well  known 
that   a   negatively    charged   body   repels   another   negatively 
charged  body  and  attracts  one  which  is  positively  charged,  it  is 
not  unreasonable  to  attribute  the  movement  of  ions  to  electrical 
attraction,  and  to  conclude  that  cathions  are  electro  positively 
charged.    Since   the   anode  is  electropositive,  the  anions  are 
thought  to  be  charged  electronegatively.    It  is  customary  to  call 
cathions  positive  ions,  and  anions  negative  ions. 

394.  The  Cause  of  the  Union  of  Ions. — Attention  has  been 
called  to  the  fact  that  in  reactions  in  solution  basic  or  metallic 
radicals  unite  only  with  acidic  or  non-metallic  radicals  (383), 
and  that  unions  of  basic  radicals  with  one  another  never  occur; 
that  is,  double  decompositions  in  water  solutions  never  give 
compounds  such  as  KNa  and  C1SO4.    We  are  now  in  position  to 
explain  these  important  facts.    We  have  learned  that  the  radicals 
of  acids,  bases,  and  salts  are  identical  with  their  ions,  and  that 


236  Introduction  to  General  Chemistry 

the  ions  are  probably  electrically  charged,  the  basic  or  metallic 
ones  being  positively,  the  acidic  or  non-metallic  negatively, 
charged.  We  can  therefore  summarize  by  stating  that  in 
reactions  only  ions  of  unlike  electric  charges  unite  with  one  another. 
The  reason  for  this  is  that  ions  with  unlike  electric  charges 
attract  each  other,  and  that  those  with  like  charges  repel  each 
other.  The  chemical  union  of  ions  is  an  electrical  phenomenon 
and  is  due  to  the  attractive  force  of  unlike  electric  charges  carried 
by  the  ions. 

395.  Colors  of  Ions. — The  student  has  doubtless  already 
observed  that,  although  most  salts  and  their  solutions  are  color- 
less, a  considerable  number  are  colored.    A  little  investigation 
will  show  that  very  dilute  solutions  of  equal  concentration  of 
all  cupric  salts  of  colorless  acids  are  of  the  same  shade  and 
intensity  of  blue  color.    This  fact  leads  us  to  believe  that  the 
blue  color  is  due  to  copper  ions,  which  is  the  only  substance  which 
all  the  solutions  have  in  common.     Moreover,  we  find  that  the 
colors  of  all  dilute  solutions  of  colored  acids,  bases,  and  salts 
can  be  ascribed  to  the  colors  of  their  ions. 

If  the  dilute  solution  of  any  acid,  base,  or  salt  is  colorless, 
like  pure  water,  we  may  conclude  that  its  positive  and  negative 
ions  are  both  colorless.  If  a  dilute  colored  solution  of  an  elec- 
trolyte has  one  colorless  ion  we  conclude  that  the  observed 
color  is  that  of  the  other  ion.  Thus  we  find  that  all  dilute 
ferrous  solutions  (173,  331)  are  pale  green  and  conclude  that 
ferrous  ion  is  pale  green.  Manganous  salts  (342)  (for  example, 
MnCl2  and  MnS04)  give  pale  pink  solutions,  therefore  positive 
Mn  ion  is  pale  pink.  On  the  other  hand,  dilute  solutions  of  all 
permanganates  (343)  are,  like  KMnO4,  deep  purple,  and  we 
conclude  that  negative  MnO4  ion  is  purple.  Similar  reasoning 
leads  us  to  conclude  that  negative  Cr04  ion  is  yellow  (345)  a*nd 
negative  Cr2O7  ion  is  orange  (345),  while  positive  Cr  ion  is  violet 
(344).  The  color  of  a  dilute  solution  is  usually  an  indication  of 
the  nature  of  one  of  its  ions. 

396.  Colors  of  Molecules.     Dissociation  of  Molecules  into 
Free  Ions. — Although  dilute  solutions  of  all  cupric  salts  are 
blue  the  solid  salts  and  also  their  concentrated  solutions  are  in 


The  Ionic  Hypothesis  237 

several  cases  of  a  different  color.  Thus  cupric  chloride,  CuCl2, 
in  solid  form  and  in  concentrated  solution  is  green,  and  cupric 
bromide,  CuBr2,  which  forms  almost  black  crystals,  gives  a 
concentrated  solution  which  is  dark  brown;  but  if  this  brown 
solution  is  sufficiently  diluted  the  color  gradually  changes  to 
blue,  finally  reaching  the  same  shade  of  color  as  that  of  any 
other  equally  dilute  cupric  solution.  A  simple  explanation 
of  these  color  changes  is  found  in  the  assumption  that  the  dark 
brown  color  is  that  of  the  molecules,  CuBr2,  while  the  blue  color 
is  due  to  Cu  ions.  From  the  fact  that  many  dilute  solutions 
of  bromides  are  colorless  we  conclude  that  Br  ions  are  colorless. 
By  following  up  this  idea  we  are  led  to  a  very  remarkable  con- 
clusion, namely,  that  molecules  of  CuBr2  exist  only  in  the  solid 
state  and  in  concentrated  solutions  but  not  to  an  appreciable 
extent  in  very  dilute  solutions.  This  is  accounted  for  if  we 
assume  as  the  concentrated  solution  is  diluted  molecules  gradually 
split  up  or  dissociate  into  ions,  thus : 

CuBr2->Cu-f2Br, 

so  that  in  a  dilute  solution  the  substance  exists  largely  as  free  Cu 
and  Br  ions.  If  we  evaporate  the  dilute  blue  solution  we  again 
obtain  a  brown  concentrated  solution  and  finally  brown  crystals. 
We  must  therefore  assume  that  the  change  is  a  reversible  one,  the 
ions  reuniting  to  form  molecules  as  the  solution  is  evaporated. 
Further  evidence  of  the  existence  of  free  ions  is  afforded  by  the 
experiments  next  to  be  considered. 


FIG.  47 

397.  The  Migration  of  Ions.— Let  us  take  advantage  of  the 
color  of  ions  to  discover  their  behavior  during  the  process  of 
electrolysis.  In  the  U-tube,  Fig.  47,  we  may  put  a  solution  of 
a  colored  electrolyte  in  the  lower  layer  and  colorless  electrolytes 


238  Introduction  to  General  Chemistry 

in  the  layers  next  to  the  electrodes.  As  colored  electrolytes  we 
may  use  copper  nitrate  or  potassium  permanganate.  When  the 
current  is  turned  on,  the  boundary  of  each  colored  electrolyte 
slowly  moves  up  into  one  of  the  colorless  layers  above  it,  just 
as  we  would  expect  if  the  colored  materal  is  the  free  ion  which 
carries  a  charge  of  electricity.  Thus  positive  copper  ion  migrates 
toward  the  negative  electrode,  and  negative  permanganate  ion 
migrates  toward  the  positive  electrode.  We  can  carry  out  an 
experiment  with  a  mixture  of  these  two  colored  salts  in  the  lower 
layer.  The  purple  layer  now  shows  on  the  side  of  the  positive 
electrode,  and  the  blue  layer  shows  on  the  side  of  the  negative 
electrode  just  as  before.  Thus  we  find  that  each  ion  migrates 
just  as  though  the  other  were  not  there;  and  this,  in  fact,  is 
just  what  we  should  expect  if  a  dilute  solution  contains  free 
ions  formed  by  dissociation  of  its  molecules. 

398.  The  Mechanism  of  Electrolysis. — We  can  now  make  a 
fairly  complete  picture  of  the  mechanism  of  the  conduction  of 
the  current  through  a  solution  and  of  the  accompanying  elec- 
trolysis. We  shall  assume  that  in  the  dilute  solution  the  dis- 
solved substance  is  partially  dissociated  into  positive  and 

negative  ions.  Fig.  48 
represents  diagrammati- 
cally  such  a  solution 
in  which  the  two  elec- 

,  ,.          , 

trodes  are  dipped,  con- 
nected with  a  pair  of 

FIG    g  dry  cells.    The  cells 

charge  the  electrodes, 

one  positively,  the  other  negatively.  The  influence  of  these 
charges  is  felt  by  the  ions  in  the  solution.  The  positive  ions 
are  attracted  by  the  negative  electrode  and  repelled  by  the 
positive  electrode  and  in  consequence  migrate  toward  the 
former.  The  negative  ions  move  in  the  opposite  direction  for 
similar  reasons. 

When  ions  come  into  contact  with  the  electrodes  of  unlike 
sign  they  give  up  their  charges  to  the  electrodes.  This  tends 
to  discharge  the  latter,  but  the  battery  keeps  them  charged  by 


Q  e  e  e 


The  Ionic  Hypothesis 


239 


continuously  sending  a  current  of  electricity  through  the  wires. 
A  more  detailed  description  of  the  mechanism  of  electrolysis, 
in  terms  of  the  newer  views  of  the  nature  of  electricity,  will  be 
given  in  chapter  xx. 

399.  Faraday's  Laws  of  Electrolysis. — As  the  result  of  care- 
ful experimental  investigation  of  the  quantities  of  substances 
liberated  during  electrolysis,  Faraday  arrived  at  the  following 
conclusions : 


FIG.  49 

1 .  The  amount  of  a  given  substance,  say  hydrogen,  set  free  by 
electrolysis  is  directly  proportional  to  the  quantity  of  electricity 
which  is  passed  through  the  solution. 

2.  The  amount  of  a  substance,  hydrogen  for  example,  which 
is  liberated  by  a  fixed  quantity  of  electricity  is  the  same,  whatever 
the  nature  of  the  solution  electrolyzed,  provided  that  this  substance 
and  nothing  else  is  liberated  at  the  given  electrode.     These  two 
statements  are  known  as  Faraday's  Laws  of  Electrolysis. 

400.  Two  Electrical  Units. — To  understand  these  laws  fully 
we  must  review  briefly  some  fundamental  facts  about  the  elec- 
trical current  so  that  we  can  appreciate  what  is  meant  by 
quantity  of  electricity.  In  the  first  place  we  know  that  if  a 
current  passes  through  a  wire  there  is  produced  around  the  wire 
a  magnetic  field.  If  we  attach  a  thread  to  the  middle  of  a 
magnetized  steel  needle  and  suspend  the  latter  above  and 
parallel  to  a  wire,  then  as  soon  as  we  cause  a  current  of  elec- 
tricity to  pass  through  the  wire  the  needle  sets  itself  at  an  angle 
to  the  former,  Fig.  49.  The  greater  the  angle  between  the 


240  Introduction  to  General  Chemistry 

needle  and  the  wire  the  stronger  is  the  magnetic  field,  and  the 
stronger  the  current  is  said  to  be.  It  is  on  this  principle  that 
instruments  are  built  to  measure  current  strength.  Of  course, 
to  measure  anything  we  must  first  adopt  some  fundamental 
unit  by  comparison  with  which  the  measurement  can  be  made. 
This  was  done  in  the  case  of  the  electric  current  on  the  basis  of 
the  strength  of  the  magnetic  field  about  a  conductor,  and  this 
unit  was  called  the  ampere.  The  ammeter 
(Fig.  50)  allows  us  to  read,  from  the  posi- 
tion occupied  by  the  needle  on  its  scale,  just 
how  many  amperes  of  current  are  passing. 
The  amperage  tells  us  the  strength  of 
the  current,  but  we  must  also  know  the 
time  during  which  the  current  is  allowed  to 


©       w       0 


FIG.  50  Pass  if  we  are  to  know  the  amount  of  elec- 

tricity delivered  at  the  terminals  of  the  con- 
ductor, say  at  two  electrodes.  //  a  current  of  one  ampere  is 
allowed  to  flow  one  second  it  is  said  to  deliver  a  unit  quantity  of 
electricity,  and  this  unit  is  called  the  coulomb. 

401.  Illustration  of  Faraday's  Laws. — The  following  facts 
will  serve  to  illustrate  the  meaning  of  Faraday's  laws.     By  the 
electrolysis  of  dilute  acids  hydrogen  is  set  free  at  the  negative 
electrode.     In   all   cases   the  passage  of  96,500  coulombs  of 
electricity  is  required  for  the  liberation  of  one  gram  of  hydrogen. 
Since  a  current  of  one  ampere  delivers  one  coulomb  per  second, 
96,500  coulombs  will  be  given  by  a  current  of  one  ampere  in 
96,500  seconds,  or  26.8  hours.     A  current  of  two  amperes  for 
the  same  length  of  time  will  liberate  2  g.  of  hydrogen,  or  one 
gram  molecular  weight  of  hydrogen  (H2) ,  which  as  we  know  has 
a  volume  of  22.4  liters  at  o°  and  76  cm. 

402.  Discussion. — It  is  not  surprising  that  if  a  one-arnpere 
current  will  liberate  i  g.  of  hydrogen  in  26 . 8  hours,  a  two-ampere 
current  will  liberate  2  g.  of  hydrogen  in  the  same  time,  for  this 
is  the  type  of  regularity  which  we  have  become  accustomed  to 
expect  in  nature.     It  is  surprising,  however,  that  the  same 
amount  of  hydrogen  is  liberated  by  the  same  amount  of  elec- 
tricity from  a  solution  of  any  dilute  acid,  and  the  fact  that  this 


T/te  Ionic  Hypothesis 


241 


is  so  must  reflect  some  regularity  in  the  phenomena  of  elec- 
trolysis, the  cause  of  which  we  have  still  to  discover. 

403.  Faraday's  Laws  of  Electrochemical  Equivalents. — Let 
us  now  turn  to  cases  of  the  liberation  by  electrolysis  of  elements 
other  than  hydrogen.  Very  careful  experimentation  has  shown 
that  by  the  passage  of  96,500  coulombs  of  electricity  through 
various  solutions  certain  weights  of  elements  are  set  free.  These 
are  given  in  Table  XV.  This  table  shows  a  most  striking  regu- 

TABLE  XV 

ELECTOCHEMICAL  EQUIVALENTS 


Element 

Gram  Atomic 
Weight 

Valence 

Weight  Liberated 
in  Grams 

Gram  Atomic 
Weight  -T-  Valence 

Hydrogen  
Silver  

1  08 

' 

I 
1  08 

I 
1  08 

Copper.  . 

64 

2 

T.2 

T.2 

Zinc 

66 

2 

7  •? 

77 

Lead 

208 

2 

IO4. 

I  O4. 

Iron 

r6 

2 

28 

28 

Aluminum  
Chlorine          .... 

27 

2C  .  C 

3 

I 

9 

7,e    r 

9 

7tr    <r 

Bromine  

80 

I 

80 

80 

Oxygen.  . 

16 

2 

8 

8 

larity:  The  weight  of  an  element  liberated  in  electrolysis  by  the 
passage  of  96,500  coulombs  of  electricity  is  equal  to  the  gram  atomic 
weight  of  that  element  divided  by  its  valence  (col.  5) .  This  weight 
is  called  the  electrochemical  equivalent  of  a  given  element  or, 
more  briefly,  its  equivalent  weight.  The  electrochemical  equiva- 
lents of  the  various  elements  are  seen  to  be  proportional  to  the 
weights  of  these  elements  which  unite  chemically  with  one 
another  when  union  is  possible;  for  example,  i  g.  of  hydrogen, 
104  g.  of  lead,  or  9  g.  of  aluminum  unite  with  35 . 5  g.  of  chlorine, 
or  8  g.  of  oxygen.  The  discovery  of  facts  such  as  those  given 
in  the  table  was  made  by  Faraday,  who  stated  his  conclusion  as 
the  Law  of  Electrochemical  Equivalents :  The  amounts  of  different 
substances  liberated  by  the  same  quantity  of  electricity  are  propor- 
tional to  their  equivalent  weights. 

404.  The  Electric  Charges  of  Ions.— The  facts  covered  by 
Faraday's  laws  allow  us  to  draw  some  interesting  and  significant 


242  Introduction  to  General  Chemistry 

conclusions  regarding  the  quantities  of  electricity  composing 
the  charges  on  single  ions.  If  the  96,500  coulombs  of  electricity 
supplied  at  the  negative  electrode  to  release  one  gram  of  hydro- 
gen ion  are  used  simply  to  neutralize  the  charge  on  one  gram  of 
that  ion,  we  may  conclude  at  once  that  the  charge  carried  by 
the  one  gram  of  hydrogen  ion  is  not  only  opposite  in  sign  but 
equal  in  amount  to  the  electricity  required.  In  general  then 
one  formula  weight  of  a  univalent  ion  must  carry  a  total  charge 
equal  to  96,500  coulombs.  One  formula  weight  of  an  ion  of  any 
valence  will  carry  one,  two,  three,  four,  etc.,  times  this  charge, 
according  to  whether  its  valence  is  one,  two,  three,  or  four,  etc. 
If  we  assume  that  one  formula  weight  of  any  one  ion  represents 
the  same  number  of  free  ions  as  a  formula  weight  of  any  other 
ion  (and  this  is  in  strict  accord  with  our  accepted  definition  of 
the  term  formula  weight),  we  come  at  once  to  the  conclusion  that 
all  univalent  ions  carry  equal  charges.  We  call  this  a  unit  charge ; 
each  bivalent  ion  carries  two  unit  charges,  each  trivalent  ion  carries 
three  unit  charges,  etc.  In  writing  the  symbols  or  formulae  of 
free  ions  it  is  customary  to  add  one  or  more  +  or  —  signs  to 
indicate  the  number  of  positive  or  negative  unit  electric  charges 
carried  by  the  ion,  for  example,  H+,  Cu++,  A1+.++,  Cl~, 

S04",  P04 . 

405.  Equilibrium  between  Molecules  and  Ions.— The  facts 
already  studied  (396),  together  with  a  great  volume  of  other 
evidence  which  we  shall  take  up  in  turn,  led  the  Swedish  chemist 
Svante  Arrhenius  to  the  conclusion  that  in  concentrated  solu- 
tions of  acids,  bases,  and  salts  a  considerable  part  of  the  dissolved 
substance  is  present  as  molecules;  but  that  as  the  solution  is 
diluted,  more  and  more  of  the  molecules  dissociate  into  free  ions, 
until  in  very  dilute  solutions  (at  least  in  many  cases)  the  disso- 
ciation is  nearly  complete.  On  the  other  hand,  when  a  dilute 
solution  is  evaporated  the  ions  undoubtedly  gradually  unite 
to  form  molecules,  until,  when  complete  dryness  is  reached,  only 
molecules  are  present.  In  any  given  solution  a  state  of  equi- 
librium exists  between  molecules  and  ions,  as  represented  in  the 
case  of  common  salt  by  the  equation 

NaCl±*Na++Cl-. 


The  Ionic  Hypothesis 


243 


At  a  definite  concentration  a  definite  proportion  of  the  salt  will 
be  present  as  ions;  this  proportion  we  call  the  fraction  ionized 
or  the  degree  of  ionization.  We  shall  next  take  up  the  important 
problem  of  determining  the  fraction  ionized  for  any  solution  of 
an  electrolyte.  Since  we  believe  that  the  current  in  a  solution 
is  carried  by  the  ions  present,  the  ability  of  a  solution  to  conduct 
a  current,  or,  briefly,  its  conductivity,  must  be  an  indication  of 
the  extent  to  which  its  molecules  are  dissociated  into  ions. 

406.  Effect  of  Dilution  on  Conductivity. — We  have  already 
learned  from  the  color  changes  produced  by  diluting  solutions 
that  ionization  is  promoted  by  dilution.  Let  us  now  consider 
the  question,  What  influence,  if  any,  does  the  volume  of  water 
in  which  a  given  quantity  of  an  acid,  base,  or  salt  is  dissolved 


HCI 


FIG.  51 

have  on  its  electrical  conductivity?  We  may  study  this  ques- 
tion experimentally  by  means  of  the  apparatus  shown  in  Fig.  5 1 . 
The  rectangular  glass  vessel  of  about  i  liter  capacity  is  provided 
with  two  large  copper  electrodes,  as  shown  in  the  figure.  The 
vessel  is  first  filled  about  three-fourths  full  of  distilled  water, 
and  the  electrical  connections  are  made.  No  appreciable  current 
passes.  Next  about  200  c.c.  of  concentrated  hydrochloric  acid 
are  introduced  below  the  water,  without  mixing,  in  such  a  way 
as  to  form  a  separate  layer.  This  may  be  done  by  the  use  of  a 
dropping  funnel,  the  stem  of  which  reaches  the  bottom  of  the 
vessel. 

The  vessel  now  contains  two  distinct  layers— a  lower  layer 
of  concentrated  hydrochloric  acid  and  an  upper  layer  of  water. 
The  galvanometer  indicates  that  a  considerable  current  is 
passing,  and  we  conclude  that  this  is  all  passing  through  the 
acid  in  the  lower  layer  and  not  through  the  upper  water  layer. 


244  Introduction  to  General  Chemistry 

If  next  we  mix  the  acid  and  water  thoroughly  and  so  dissolve 
the  acid  in  a  much  larger  volume  of  water,  we  note  that  a  large 
increase  in  the  current  takes  place.  This  leads  us  to  conclude 
that  the  conductivity  of  the  hydrochloric  acid  present  is  greater 
when  it  is  dissolved  in  the  larger  volume  of  water.  We  may  now 
ask,  however,  whether  there  is  a  limit  to  the  increase  in  conduc- 
tivity when  a  given  amount  of  acid  is  dissolved  in  larger  and 
larger  volumes  of  water,  the  conductivity  being  measured  under 
such  conditions  that  the  solution  is  all  contained  between  parallel 
plates  at  a  fixed  distance  apart. 

If,  in  the  experiment  described,  the  vessel  were  much  deeper, 
but  otherwise  the  same,  and  the  electrodes  extended  all  the  way 
up  the  sides  as  before,  it  would  be  found  that  a  given  amount 
of  hydrochloric  acid  diluted  with  double  the  amount  of  water 
used  in  the  first  experiment  would  show  appreciably  greater 
conductivity  than  in  the  first  case.  Or  if  the  acid  were  diluted 
with  three,  or  four,  or  still  more  times  as  much  water,  greater 
and  greater  conductivity  would  have  been  observed;  but  with 
increasing  dilution  the  increase  in  conductivity  would  become 
smaller  each  time  more  water  was  added,  so  that  finally  a  maxi- 
mum conductivity  would  be  reached.  Beyond  this  limit  further 
dilution  would  cause  no  increase  in  conductivity. 

These  same  experiments  could  be  repeated  with  many  other 
electrolytes  with  similar  results. 

407.  Definition  of  Molecular  Conductivity. — If  one  formula 
weight  (called  also  one  gram  molecular  weight)  of  an  acid,  base, 
or  salt  is  contained  in  a  solution  which  is  wholly  included  between 
two  parallel  electrodes  i  cm.  apart,  we  call  the  electrical  con- 
ductivity of  this  solution  its  molecular  conductivity.  To  find 
the  molecular  conductivity  experimentally  we  measure  its 
electrical  resistance  in  ohms.  The  reciprocal  of  the  resistance 
so  found  is  by  definition  the  molecular  conductivity.  The  con- 
clusions of  the  paragraph  on  the  effect  of  dilution  on  conductivity 
may  now  be  summarized  as  follows:  The  molecular  conductivity 
of  every  electrolyte  increases  as  its  solution  is  diluted  and  finally 
attains  a  maximum  which  has  a  definite  numerical  value  for  each 
substance  (the  temperature  being  fixed).  Table  XVI  shows  the 


The  Ionic  Hypothesis  245 

change  of  molecular  conductivity  of  hydrochloric  acid  as  the 
volume  in  which  one  formula  weight  of  acid  is  contained  is 
increased. 

TABLE  XVI 

THE  MOLECULAR  CONDUCTIVITY  OF  HYDROCHLORIC 
ACID  AT  l8°  (NOYES  AND  COOPER) 

Volume  of  Solution  Molecujar 

in  Liters  Conductivity 

10 351 

12.5 353 

100 368 

5oo 373 

2,000 375 

Maximum 379 

408.  Determination  of  the  Degree  of  lonization — When  a 
solution  of  a  substance  is  so  dilute  that  it  has  its  maximum  molec- 
ular conductivity,  it  is  assumed  that  all  of  its  molecules  have 
dissociated  into  ions.  In  a  more  concentrated  solution,  for  which 
the  molecular  conductivity  is  less  than  the  maximum,  the  frac- 
tion which  its  observed  molecular  conductivity  forms  of  its 
maximum  molecular  conductivity  is  consequently  equal  to  the 
fraction  which  the  number  of  ions  present  in  that  particular 
solution  form  of  the  total  number  of  ions  in  the  completely 
dissociated  (completely  ionized)  solution  of  the  same  quantity 
of  that  substance.  This  fraction  is  therefore  the  fraction  ionized 
or  the  degree  of  ionization.  Thus  for  decinormal  hydrochloric 
acid  the  degree  of  ionization  is  351-1-379  =  93  per  cent. 

This  method  of  determining  the  degree  of  ionization  was 
proposed  by  Arrhenius  in  the  year  1887.  His  reasoning  ran 
thus:  The  passage  of  a  current  through  a  solution  is  accom- 
plished by  the  migration  of  positive  ions  in  one  direction  and 
negative  ions  in  the  other.  These  transport  electricity  through 
the  solution  between  the  electrodes.  Since  the  molecular  con- 
ductivity of  a  substance  is  the  measure  of  its  rate  of  transporting 
electricity,  it  is  plain  that  the  molecular  conductivity  will  depend 
on  the  number  of  ions  present,  the  charge  on  each,  and  the  velocity 
of  migration.  Now  under  the  conditions  used  in  measuring 
resistance,  and  therefore  also  of  measuring  molecular  con- 
ductivity, the  velocity  of  migration  of  its  ions  will  be  the  same 


246  Introduction  to  General  Chemistry 

for  all  concentrations  of  solutions  of  a  given  substance  (except 
for  very  concentrated  solutions).  The  charges  of  the  individual 
ions  of  a  given  substance  are  also  the  same,  whether  the  solution 
is  dilute  or  concentrated.  Therefore  the  molecular  conductivities 
•of  solutions  of  a  given  substance  are  directly  proportional  to  the 
numbers  of  ions  present.  Consequently  the  ratio  of  the  molecu- 
lar conductivity  for  a  given  concentration  to  the  maximum  mo- 
lecular conductivity  for  this  substance  is  the  fraction  ionized,  since 
it  is  assumed  that  a  very  dilute  solution  having  maximum  molec- 
ular conductivity  is  completely  ionized. 

409.  Results  of  Determination. — The  degree  of  ionization  of 
some  common  electrolytes  is  shown  in  Table  XVII. 

410.  Discussion  of  Table  XVII. — A  study  of  the  table  leads 
to  the  very  important  generalization  that  in  solutions  of  most 
salts  a  large  percentage  of  the  substance  is  in  the  form  of  ions;  in 
consequence,  we  say  that  such  solutions  are  highly  ionized.    It 
also  appears  that  dilute  solutions  of  hydrochloric  and  nitric  acid 
are  even  more  highly  ionized  than  salt  solutions  of  like  concentra- 
tion.   On  the  other  hand,  decinormal  acetic  acid  is  only  i .  3  per 
cent  ionized,  while  the  degree  of  ionization  of  decinormal  carbonic 
acid  is  very  much  less,  namely  0.17  per  cent.     In  general,  the 
extent  to  which  acids  are  ionized,  in  solutions  of  equal  concen- 
tration, varies  enormously.    Bases  also  differ  greatly  in  their  de- 
grees of  ionization.     For  example,  decinormal  sodium  hydroxide 
is  ionized  90  per  cent,  while  the  same  concentration  of  ammonium 
hydroxide  is  only  i .  3  per  cent  ionized.    We  have  already  learned 
that  every  substance  is  more  highly  ionized  in  dilute  than  in 
more  concentrated  solutions.    The  percentage  of  ionization  of 
a  substance  as  shown  in  the  table  applies  only  to  the  indicated 
concentration  and  temperature. 

411.  Resume  of  the  Ionic  Hypothesis. — We  have  already 
developed  enough  of  the  ionic  hypothesis  to  go  far  into  the  under- 
standing of  double  decomposition  reactions.    Let  us  therefore 
review  in  brief  the  ideas  already  brought  out,  and  then,  after 
a  short  critical  survey  of  the  fundamental  assumptions,  proceed 
in  chapters  xviii  and  xix  to  the  application  of  the  hypothesis  to 
practical  examples. 


The  Ionic  Hypothesis 


247 


According  to  the  ionic  hypothesis,  as  soon  as  an  acid  base  or 
salt  is  dissolved  in  water  it  is  immediately  dissociated  to  some 
extent  into  ions  which  prove  to  be  the  parts  of  those  substances 
which  we  have  found  active  in  double  decomposition.  The  basic 

TABLE  XVII 

VALUES  OF  THE  DEGREE  OF  IONIZATION  OF  SOME  COMMON  ELECTROLYTES  IN 

WATER  SOLUTION  AT  18° 
(Degree  of  ionization  at  the  normality  indicated  at  the  head  of  the  column) 


0    01 

O.OS 

O.I 

I  .0 

SALTS: 
NaCl 

o  04. 

o  88 

o  8s 

O    74 

KC1  

04- 

88 

86 

74 

KBr  

04 

88 

86 

KI  

•  04 

80 

87 

73 

NaNO3  

.02 

87 

83 

.66 

KNO3  

04 

.87 

.82 

.61 

AgNO, 

Q2 

86 

81 

62 

KC1O3 

03 

87 

8* 

BaClj 

88 

80 

* 

76 

64 

CaCU 

88 

80 

76 

66 

MgCl2 

.88 

.80 

77 

67 

PbCl2  

.81 

.6* 

Sr(NO3)2   

.87 

,  77 

•  72 

•  *>I 

Ba(NO3)3 

86 

74. 

68 

K2SO4  

.87 

•77 

.72 

•59 

Ag2SO4 

84 

MeSCX 

67 

PI 

At 

.  3"? 

ZnSO4 

63 

46 

41 

•  ^i 

CuSO4         .       ... 

•  63 

.46 

.40 

o.  31 

BASES: 
NaOH 

06 

Q-2 

oo 

Ba(OH)2 

Q2 

86 

81 

NH4OH 

O4 

Ol7 

01^ 

ACIDS: 
HC1  

.07 

•  04 

.0? 

HNO3 

O7 

O4 

Q7 

HC2H3O2 

O42 

O2O 

OH 

H3PO4  _  H++H2PO4~ 

60 

^6 

20 

H2SO4=2H++SO4  —  ".... 
H2CO3  —  H++HCO3- 

.64 

o  005 

.38 
O   OO2 

•31 

o  0017 



part  carries  a  positive  charge  and  the  acidic  part  a  negative 
charge.  The  charge  carried  by  any  univalent  ion  is  called  a 
unit  charge;  ions  having  greater  valence  carry  as  many  unit 
charges  as  they  have  valence.  Since  the  solution  of  any  elec- 
trolyte is  always  electrically  neutral,  the  total  quantity  of  posi- 
tive electricity  carried  by  the  positive  ions  equals  the  total 


248  Introduction  to  General  Chemistry 

quantity  of  negative  electricity  carried  by  the  negative  ions. 
The  more  dilute  the  solution  the  greater  is- the  proportion  of  the 
electrolyte  transformed  into  ions,  or,  in  other  words,  the  greater 
is  its  degree  of  ionization. 

When  we  put  two  electrodes,  connected  with  dry  cells  or 
other  source  of  current,  into  a  solution  of  an  electrolyte,  the 
current  is  found  to  flow  in  the  outside  circuit,  because  of  the 
discharging  of  the  charge  on  the  electrodes,  due  to  the  arrival 
at  their  surface  of  oppositely  charged  ions  from  the  solution. 
The  ions  in  the  solution  move  up  to  the  plates  because  each  ion 
is  attracted  toward  one  plate  and  repelled  from  the  other,  owing 
to  the  fact  that  it  also  carries  a  charge  of  electricity.  The  sign 
of  the  charge  on  the  ion  determines  the  direction  of  the  latter's 
movement.  On  coming  in  contact  with  the  electrodes  the  ions 
become  discharged  by  having  their  charges  electrically  neutral- 
ized by  equal  amounts  of  the  opposite  kind  of  electricity  furnished 
by  the  electrode.  Metallic  ions  after  discharge  are  either  de- 
posited as  metallic  platings  on  the  cathode  or  react  with  water 
to  form  hydroxides  and  hydrogen.  Non-metallic  ions  such  as 
hydrogen,  oxygen,  and  chlorine  are  released  as  single  atoms  which 
then  unite  to  form  diatomic  molecules  of  the  gases  H2,  O2,  C12, 
etc.  Nitrate  and  sulfate  ions  never  remain  discharged  at  the 
electrodes,  but  instead  we  find  there  the  products  of  their 
reactions  with  water — nitric  acid  and  oxygen,  and  sulfuric  acid 
and  oxygen,  respectively. 

If  a  given  quantity  of  electrolyte  is  kept  between  plates 
which  are  parallel  to  each  other  and  carry  a  constant  charge  per 
unit  area,  at  different  dilutions  the  conductivity  of  this  elec- 
trolyte will  vary  in  proportion  to  the  number  of  free  ions  present. 
As  a  consequence  the  proportion  of  an  electrolyte  which  has 
been  transformed  into  ions  at  any  dilution  can  be  determined 
by  dividing  the  conductivity  at  the  dilution  in  question  by  the 
maximum  conductivity  found  after  continuous  dilution  of  this 
same  quantity  of  electrolyte,  provided  the  two  measurements 
are  made  in  the  manner  described  (408).  Values  so  determined 
show  us  that  as  a  rule  all  salts  are  highly  ionized  substances, 
but  acids  and  bases  have  very  different  degrees  of  ionization, 


The  Ionic  Hypothesis  249 

some  being  even  more  highly  ionized  than  salts,  but  others 
being  very  little  ionized  indeed. 

412.  Criticisms  of  the  Ionic  Hypothesis. — The  idea  that  ions 
exist  in  solution  as  independent  chemical  substances  has  come 
to  be  known  as  the  Ionic  Hypothesis.  It  will  be  surprising  if 
the  student  who  learns  of  this  hypothesis  for  the  first  time  and 
thinks  critically  about  the  matter  is  not  ready  to  offer  at  once 
several  good  reasons  for  doubting  the  truth  of  the  conclusions. 
In  the  first  place,  the  hypothesis  seems  to  assume  that  in  a  solu- 
tion of  common  salt,  for  example,  a  large  part  of  its  elements 
are  present  in  a  free  state.  Now  the  student  who  knows  any- 
thing of  the  properties  of  metallic  sodium  and  of  chlorine  gas 
will  find  it  hard  to  believe  that  either  of  these  elements  can  be 
present  in  a  salt  solution;  because  sodium  reacts  violently  with 
water,  forming  sodium  hydroxide  and  hydrogen,  and  chlorine 
has  a  horrible  smell  and  a  yellow  color.  Plainly  there  is  some- 
thing incompatible  with  the  obvious  facts  in  the  statement  that  a 
solution  of  salt  contains  free  sodium  and  free  chlorine. 

A  closer  study  of  the  hypothesis  shows,  however,  that  it  is 
not  assumed  that  the  elements  sodium  and  chlorine  are  present 
as  ordinary  atoms,  for  each  atom  is  said  to  be  electrically  charged. 
Those  who  uphold  the  hypothesis  will  point  out  that  a  charged 
brass  ball  has  very  different  properties  from  the  same  ball  if 
uncharged.  True,  say  the  critics,  but  even  a  charged  brass  ball 
is  still  a  brass  ball;  to  which  the  opponents  reply  that  the 
quantity  of  electricity  on  the  ball  is  a  matter  of  enormous 
importance. 

If  then  the  ions  are  so  highly  charged,  why  do  the  positive 
ions  not  unite  with  and  so  electrically  discharge  the  negative 
ions,  since  the  solution  is  a  conductor?  It  may  be  said  in  reply 
that  it  is  assumed  that  ions  of  unlike  sign  are  constantly  uniting, 
at  a  rate  just  equal  to  the  rate  of  dissociation,  with  the  result 
that  a  state  of  equilibrium  is  produced. 

In  spite  of  the  foregoing  criticisms  and  many  others  the  ionic 
hypothesis  with  all  its  apparent  inconsistencies  has  proved  itself 
highly  useful  in  explaining  and  correlating  many  facts  and 
phenomena. 


250  Introduction  to  General  Chemistry 

Before  passing  final  judgment  on  this  remarkable  hypothesis 
it  will  be  better  to  consider  its  further  applications  and  then, 
in  chapter  xx,  to  take  up  the  matter  again  in  the  light  of  newer 
discoveries,  which  have  led  to  essential  modifications  of  the 
views  as  originally  proposed  by  Arrhenius. 

Finally  it  may  be  urged  that  Arrhenius  himself  was  not 
certain  of  the  truth  of  his  theory  until  he  became  acquainted 
with  the  wonderful  work  of  Van't  Hoff  on  the  so-called  osmotic 
pressures  of  dissolved  substances  (chap,  xxvii).  This  work 
will  be  discussed  as  soon  as  we  have  progressed  far  enough  to 
understand  and  interpret  the  experiments  which  we  must  then 
study. 


CHAPTER  XVIII 
APPLICATIONS  OF  THE  IONIC  HYPOTHESIS 

413.  Double  Decomposition.  —  In  the  foregoing  chapter  it 
was  pointed  out  that  the  probable  cause  of  the  union  of  two 
unlike  ions  is  the  attraction  of  their  unlike  electric  charges.     In 
general,  every  kind  of  positive  ion  can  unite  with  any  kind  of 
negative  ion.     Therefore,  if  any  two  electrolytes  (provided  they 
have  no  ion  in  common)  are  mixed  in  solution,  at  least  some 
double  decomposition  must  take  place,   simply  because  new 
combinations  of  positive  and  negative  ions  are  made  possible. 
Let  us  first  consider  the  important  case  in  which  the  two  starting 
materials,  as  well  as  the  two  products  of  the  reaction,  are  easily 
soluble  and  highly  ionized. 

414.  Class  I.    Equilibrium  between  Four  Easily  Soluble  and 
Highly  Ionized  Electrolytes.  —  If  dilute  solutions  of  two  imaginary 
electrolytes  A  B  and  CD,  which  ionize  thus 

AB±»A++B-,  I 

CD^C++D~,  II 

are  mixed,  we  may  predict,  without  knowing  anything  more 
about  these  substances,  that  the  following  reactions  are  possible, 


III 
C++B~^CB,  IV 

and  that  a  double  decomposition  reaction, 
AB+CD^AD+CB, 

will  take  place  to  a  greater  or  less  extent.  Since  all  four  of  the 
substances  AB,  CD,  AD,  and  CB  are  assumed  to  be  highly 
ionized,  it  is  plain  that  the  mixed  solution  will  contain  chiefly  the 
four  kinds  of  ions,  A+,  B~,  C+,  and  D~,  and  relatively  few 
molecules.  Since  each  of  the  four  kinds  of  molecules  present 
must  be  in  equilibrium  with  its  own  two  kinds  of  ions,  the  four 

251 


252  Introduction  to  General  Chemistry 

equilibrium  reactions  (marked  I,  II,  III,  and  IV)  must  be 
interrelated  in  the  manner  shown  by  the  following  arrangement 
of  equations  I,  II,  III,  and  IV: 

III     IV 

AB±*A++B-  I 

CD±>D~+C+  II 

It       It 
AD    CB 

Equations  I  and  II  read  horizontally,  while  III  and  IV  read 
vertically.  We  may  call  this  the  compound  equation  of  the 

reaction 

AB+CD^AD+CB. 

The  compound  equation  shows  that  the  four  molecular  sub- 
stances are  in  equilibrium  with  each  other  because  each  molec- 
ular substance  is  in  direct  equilibrium  with  its  own  pair  of  ions. 
Now  if  all  of  the  four  molecular  substances  are  assumed  to  have 
exactly  equal  tendencies  to  ionize,  then  we  must  conclude  that 
for  the  condition  of  equilibrium  equal  numbers  of  the  four  kinds 
of  molecules  will  be  present,  if  we  have  taken  equivalent  amounts 
of  substances.  We  may  summarize  Class  I  as  follows:  If  both 
starting  substances  and  both  products  of  a  double  decomposition 
reaction  AB+CD^  AD+CB  are  easily  soluble  and  highly 
and  equally  ionizable,  an  equilibrium  mixture  will  result  in 
which  (i)  most  of  the  dissolved  material  is  present  as  free  ions, 
(2)  little  of  the  material  is  present  as  molecules,  and  (3)  if 
equivalent  amounts  are  taken  the  four  kinds  of  molecules  are 
present  in  equal  numbers. 

415.  An  Example  of  Class  I.    The  Reaction  between  Ferric 
Chloride  and  Ammonium  Sulfocyanate. — The  reaction 
FeCl3+3NH4NCS^Fe(NCS)3+3NH4Cl, 

studied  earlier  (280),  is  a  good  illustration  of  Class  I,  since  all 
four  salts  are  easily  soluble  and  highly  ionized.  It  was  shown 
by  experiment  that  this  reaction  does  not  take  place  completely, 
but  that  it  reaches  equilibrium  while  there  is  still  much  of  the 
material  not  converted  into  ferric  sulfocyanate  and  ammonium 
chloride. 


Applications  of  the  Ionic  Hypothesis  253 

416.  Other   Examples    of    Class   I.  —  Mimerous   additional 
examples  of  Class  I  might  be  given.     The  following  will  serve 
as  illustrations  : 

KCl+NaNO3^KNO3+NaCl, 

Na2SO4+  2KNO3  ±5  K2SO4+  2NaNO3, 

K2CO3+Na2S04^K2SO4+Na2CO3, 

NaNO3+HCl^NaCl+HNO3. 

In  each  case  the  mixed  solution  contains  largely  the  four  kinds 
of  ions,  together  with  small  proportions  of  the  four  kinds  of 
molecules  in  approximately  equivalent  amounts.  In  Class  I 
the  two  substances  taken  react  only  partially,  and  there- 
fore the  reaction  is  always  incomplete. 

417.  A  Graphic  Method  of  Representing  Degrees  of 
lonization.  —  An  acid,  base,  or  salt,  not  in  solution,  exists    FlG  52 
wholly  in  the  form  of  molecules  (no  ions  are  present). 

We  may  represent  such  an  un-ionized  substance  by  a  cross- 
hatched  circle,  Fig.  52.  When  this  substance,  whose  formula  we 
may  call  AB,  is  dissolved  in  water  it  will  partially  ionize,  thus: 


This  condition  is  represented  by  Fig.  53.  Let  us  suppose  that 
the  solution  is  80  per  cent  ionized;  then  20  per  cent  is  present  as 
un-ionized  molecules.  In  Fig.  53  the  left-hand  circle  has  a 

cross-hatched  sector  which,  is 
'\AB        /<|iilM  Just  20  Per  cent  °^  the  area 


of  the  whole  circle.    This  will 
represent    the   fact   that    20 
FIG.  53  per  cent  of  the  substance  is 

present   as  un-ionized  mole- 

cules. The  middle  circle,  of  which  80  per  cent  is  shaded  with 
vertical  lines,  will  represent  the  fact  that  80  per  cent  of  the 
total  A  radical  is  in  the  form  of  free  positive  ions.  In  similar 
fashion  the  right-hand  circle  shows  that  80  per  cent  of  the 
total  B  radical  is  in  the  form  of  free  negative  ions.  Further- 
more, if  we  take  the  area  of  the  circle,  Fig.  52,  as  propor- 
tional to  the  whole  number  of  molecules  in  one  formula  weight 
of  the  substance  before  it  is  dissolved,  then  the  area  of  the 


254 


Introduction  to  General  Chemistry 


CuSO* 


sor- 


cross-hatched  sector  of  the  left-hand  circle  of  Fig.  53  will  be  pro- 
portional to  the  number  of  un-ionized  molecules  in  one  formula 
weight  of  the  dissolved  substance.  Since  each  AB  molecule, 

when  it  ionizes,  gives  one  A  + 
ion  and  one  B~  ion,  the  areas 
of  the  shaded  portions  of  the 
middle  and  right-hand  circles 
will  be  directly  proportional  to 
the  numbers  of  A+  and  B~ 
ions  respectively.  By'  means 
of  a  figure  like  Fig.  53  the 
relative  concentrations  of  ions 
and  molecules  of  a  dissolved 
electrolyte  can  be  seen  at  a 
glance.  By  way  of  further 
illustration  the  condition  of 
normal,  one-  tenth-normal,  and  one  one-hundredth-normal  copper 
sulfate  solution  is  shown  in  Fig.  54. 

418.  Graphic  Representation  of  Class  I.  —  Let  us  now  turn 
to  the  graphic  representation  of  a  double  decomposition  reaction 
of  the  type  just  studied  under  Class  I,  where  all  four  substances 
concerned  are  easily  soluble  and  highly  ionized.  We  again  repre- 
sent the  reaction  by 


so«— 


FIG.  54 


AB 


Figure  55  shows  the  condition  of  solutions  of   AB  and  CD 
before  they  are  mixed,  on  the  supposition  that  each  is  85  per  cent 
ionized  in  N/  10  solution.    When 
equal  volumes  of  the  two  N/io 
solutions  are  mixed,  the  reaction 
represented    by  the  compound 
equation 


FIG.  55 


it       It 
AD    CB 


takes  place  and  very  quickly  reaches  the  condition  of  equilib- 
rium shown  graphically  in  Fig.  56,  in  which  the  proportions  of 


Applications  of  the  Ionic  Hypothesis  255 

molecules  and  ions  have  been  calculated  on  the  additional 
assumption  that  AD  and  CB  both  have  the  same  tendency  to 
ionize  as  have  A  B  and  CD  (when  each  is  separately  dissolved 
in  water).  Comparison  of  Figs.  55  and  56  shows  us  that  the 
areas  representing  the  numbers 
of  molecules  and  ions  of  the 
materials  taken  are  not  greatly 
changed  as  the  result  of  the 
mixing.  Consequently  we  say 
that  the  reaction  is  incomplete. 
All  examples  of  Class  I  would  be 
represented  by  similar  graphs. 

419.  A    Second    Type    of 
Double    Decomposition:    Class 

II.  —  Class  II  will  comprise  FIG.  56 

double  decomposition  reactions 

in  which  two  easily  soluble  and  highly  ionized  substances  give 

two  easily  soluble  products,  one  of  which  is  highly  ionized,  the 

other  little  ionized.     The  simplest  example  of  Class  II  is  found 

in  a  neutralization  reaction  such  as 

HCl+NaOH^NaCl+HA 

since  all  the  substances  except  the  water  are  highly  ionized. 

420.  The  lonization  of  Water.  —  The  ionization  of  water  may 
be  determined  from  conductivity  measurements,  for  though  it  is 
a  very  much  poorer  conductor  of  the  current  than  is  a  salt  solu- 
tion, still,  as  we  have  already  said,  it  conducts  much  better 
than  glass  or  hard  rubber.    According  to  the  ionic  theory  it  is 
ionized  thus: 


In  one  liter  of  pure  water  there  is  present  about  one  ten-millionth 
of  a  gram  of  ionic  hydrogen  and  the  equivalent  amount  of 
hydroxyl. 

If  then  we  attempt  to  represent  the  proportion  of  ions  in  pure 
water  by  a  graphic  scheme,  a  single  dot  in  the  center  of  aa  other- 
wise empty  circle  would  have  too  large  an  area'  to  represent 
correctly  the  proportion  of  ions  present  if  the  rest  of  the  circle 


256  Introduction  to  General  Chemistry 

represented  the  molecules  of  water.  In  cases  of  this  kind  we 
shall  use  a  single  dotted  radius  to  indicate  that  the  number  of 
ions  is  too  small  to  be  accurately  represented.  The  graph  of 

water  will  then  be  that  shown  in 

/'      %\H*      /"     'N0"-      T,. 

~   ;    —  j    ,;    —  :     Fig.  57. 

V._y     \__*'  That    there    are   so   few   ions 

FlG  present  in  a  liter  of  pure  water 

means  that  the  ionic  equilibrium 

is  established  only  when  all  but  a  minute  fraction  of  the  total 
material  is  in  the  form  of  water  molecules.  Accordingly,  when 
hydroxyl  and  hydrogen  ions  are  brought  together  in  solution 
we  must  expect  them  to  combine  almost  completely  to  form 
molecules. 

421.  Neutralization.  —  If  we  mix  equivalent  amounts  of 
solutions  of  HC1  and  NaOH  the  resulting  reaction  may  be 
represented  as  follows: 


NaOH±5OH-+Na+ 

it          It- 
H20    NaCl. 

The  H+  and  OH~  ions  present  unite  almost  completely  to  form 
molecules  of  H2O.  The  removal  of  H+  ions  causes  a  shift  of  the 
reaction 


to  the  right,  and  as  the  H+  ions  produced  in  this  way  are  almost 
immediately  taken  up  by  new  OH~  ions  formed  by  a  shift  to  the 
right  of  the  reaction 

NaOH±5OH-+Na+, 

the  final  result  is  the  practically  complete  dissociation  of  both 
HC1  and  NaOH  molecules  and  therefore  the  disappearance  of 
these  substances.  Molecules  of  H20,  once  formed,  dissociate 
very  little  into  H+  and  OH~  ions,  and  so  the  final  equilibrium 
solution  will  contain  no  more  free  H+  and  OH~  ions  than  an  equal 
volume  of  pure  water.  The  Na+  and  Cl~  ions  unite  partially  to 
form  molecules 

Na++Cl-±»NaCl, 


Applications  of  the  Ionic  Hypothesis 


257 


but  this  reaction  does  not  proceed  far  in  dilute  solution,  as  com- 
mon salt  is  a  highly  ionized  substance.  In  fact,  the  solution 
resulting  from  the  neutralization  of  HC1  by  NaOH  is  exactly 
the  same  as,  and  differs  in  no  way  from,  that  made  by  dissolving 
common  salt  in  water  to  produce  a  solution  of  equal  concentra- 
tion. All  the  facts  just  stated  are  shown  by  a  comparison  of  the 
two  graphs,  Figs.  58  and  59.  Thus  it  can  be  seen  (Fig.  59)  that 
the  areas  representing  the  numbers  of  molecules  of  HC1  and 
NaOH  respectively  have  been  reduced  to  negligible  dimensions; 
the  same  is  also  true  of  the  areas  for  H+  and  OH~  ions.  But 


HCI 


NaOH 


FIG.  58 


Fie.  59 


the  Cl  and  Na+  ion  areas  are  not  greatly  changed  in  the  second 
graph.  Compared  with  these  areas,  that  representing  NaCl  mole- 
cules is  small.  The  circle  representing  the  number  of  molecules 
of  H2O  is  completely  shaded,  thus  showing  that  the  yield  of 
molecular  H2O  is  practically  100  per  cent. 

422.  The  Simplified  Equation  of  Neutralization. — To  sum 
up  the  matter,  it  may  be  said  that  acids  and  bases  neutralize 
one  another  because  of  the  tendency  of  H+  and  OH~  ions  to  unite 
almost  completely  to  form  water.  This  almost  complete  union 
of  H+  and  OH~  ions  takes  place  because  H2O  is  but  very  slightly 
ionized.  In  a  very  dilute  solution,  where  the  acid  and  base 
taken  are  almost  completely  ionized  at  the  instant  of  mixing, 
the  principal  change  that  takes  place  is  the  union  of  H+  and 
OH~  ions  to  form  H2O  molecules,  since  in  the  very  dilute  solu- 
tion the  Na+  and  Cl"  ions  remain  largely  uncombined.  We  may 


258  Introduction  to  General  Chemistry 

therefore  write  as  the  simplified  equation  of  neutralization  in 
dilute  solution 


423.  Experimental  Confirmation  of  the  Theory  of  Neutraliza- 
tion. —  The  process  of  neutralization  can  be  followed  experi- 
mentally with  the  help  of  an  apparatus  somewhat  like  that  shown 
in  Fig.  51  (406);  but  having  a  small  electric  lamp  in  the  place 
of  the  galvanometer.     The  solution  layers  in  the  cell  shown 
in  the  figure  are  made  by  first  putting  into  the  cell  a  layer  of 
one-tenth-normal  hydrochloric  acid,  and  then  allowing  an  equal 
layer  of  sodium  hydroxide  to  run  under  this  first  layer  by  intro- 
ducing it  at  the  bottom  of  the  cell  through  a  dropping  funnel. 
As  represented  in  the  figure  the  two  parallel  electrodes  are  in 
contact  with  the  two  layers,  which  can  be  seen  very  nicely  if  a 
little  litmus  is  put  into  the  acid  and  base  respectively  before  the 
layers  are  made.    If  now  the  key  is  closed  the  current  flows 
through  both  layers,  and  the  lamp  glows.    Hydrogen  and  sodium 
ions  are  arriving  at  one  electrode,  and  chlorine  and  hydroxyl 
ions  are  arriving  at  the  other.    Of  these  ions  the  hydrogen  and 
hydroxyl  travel  much  more  rapidly  under  the  attraction  from 
a  given  charge  per  unit  area  of  the  electrode,  and  so  they  are 
neutralizing  their  charges  on  the  plates  more  quickly  than  are 
the  other  ions.    As  a  result  most  of  the  current  passing  in  the 
outside  circuit  is  due  to  their  discharge.     If  the  two  layers  of 
acid  and  base  are  next  mixed,  the  lamp  no  longer  glows.     Half 
of  the  carriers  of  the  current  and  the  most  efficient  ones  have 
been  used  to  form  water  molecules,  and  in  the  cell  there  remains 
only  the  slow-moving  sodium  ions  and  chlorine  ions.     If  the 
acid  and  base  in  the  respective  layers  were  not  quite  equivalent 
in  amount,  a  slight  excess  of  one  or  the  other  will  be  shown  by 
the  litmus  color,  but  the  important  part  of  the  experiment,  the 
serious  loss  of  ions,  will  still  be  unmistakable  from  the  great 
decrease  in  the  conductivity  of  the  solution  between  the  plates. 

424.  A  Second  Example  of  Class  II:    Action  of  HC1  on 
Sodium  Acetate.  —  It  will  be  recalled  that  acetic  acid,  HC2H3O2, 
neutralizes  NaOH,  forming  sodium  acetate,  thus: 


Applications  of  the  Ionic  Hypothesis  259 

Acetic  acid  is  a  monobasic  acid,  only  one  of  the  four  hydrogen 
atoms  of  each  molecule  being  ionizable: 

HC2H302^H++C2H302-. 

This  acid  is  but  little  ionized  in  normal  solution,  the  degree  of 
ionization  being  only  0.4  per  cent.  On  the  other  hand,  solu- 
tions of  its  salts,  like  NaC2H302,  are  highly  ionized: 

NaC2H3O2^Na++C2H3O2-. 

If  we  mix  equivalent  amounts  of  HC1  and  NaC2H3O2  in 
solution  we  cannot  see  that  any  chemical  change  occurs;  but 
that  a  reaction  has  occurred  we  may  show  convincingly  with  the 
help  of  the  electrolytic  cell,  which  is  used  to  discover  the  change 
in  conductivity  during  neutralization.  In  the  lower  layer  this 
time  we  shall  have  sodium  acetate  and  in  the  upper  hydrochloric 
acid.  As  before,  the  lamp  glows  —  both  solutions  are  good  con- 
ductors; the  first  by  means  of  Na+  and  C2H302~  ions,  the 
second  by  means  of  H+  and  Cl~  ions.  When  we  mix  the  two 
layers  the  decrease  in  brightness  of  the  lamp  shows  that  the 
conductivity  has  dropped  off  greatly,  thus  proving  that  many  of 
the  ions  have  been  changed  into  non-conducting  molecules. 
The  compound  equation  is 


NaC2H302±5C2H302--!-Na+ 

It  It 

HC2H3O2    NaCl 

The  graphs  are  shown  in  Figs.  60  and  61.  Since  of  the  four  sub- 
stances concerned  all  but  the  acetic  acid  are  highly  ionized,  while 
the  latter  is  but  little  ionized,  the  reaction  falls  under  Class  II. 
When  HC1  and  NaC2H3O2  solutions  are  mixed,  the  H+  and 
C2H3O2~  ions  will  unite  far  more  completely  than  will  any  other 
pair  of  ions,  and  at  the  same  time  the  molecules  of  HC1  and 
NaC2H3O2  will  continue  to  ionize  until  but  very  few  remain 
(Fig.  61).  Also  Na+  and  Cl~  ions  will  unite  partially  to  form 
molecules  of  NaCl.  Therefore  the  equilibrium  mixture  will 
contain  largely  free  acetic  acid,  for  the  most  part  un-ionized, 


260 


Introduction  to  General  Chemistry 


together  with  common  salt  and  its  ions.     Very  little  HC1  and 
NaC2H302  will  be  present. 

425.  Comparison  of  the  First  and  Second  Examples  of  Class 
n. — Fig.  60  shows  the  conditions  of  the  solutions  of  hydrochloric 
acid  and  of  sodium  acetate  before  they  are  mixed,  while  Fig.  61 
shows  the  condition  of  the  equilibrium  mixture.  These  figures 
are  almost  a  reproduction  of  Figs.  58  and  59,  representing  neutral- 
ization. In  place  of  NaOH  we  have  in  the  second  case  NaC2H3O2, 
which  is  also  highly  ionized;  and  in  place  of  water  we  have 
HC2H3O2,  which,  like  water,  is  but  little  ionized.  In  Fig.  61  the 
circle  representing  molecular  acetic  acid  is  nearly  completely 


FIG.  60 


FIG.  61 


cross-hatched,  showing  that  the  yield  of  this  substance  is  nearly 
100  per  cent.  Although  the  reactions  represented  by  Figs.  59 
and  6 1  are  so  nearly  alike,  there  is  a  small  difference  due  to  the 
fact  that  acetic  acid  is  ionized  more  than  water.  In  consequence 
the  formation  of  molecular  acetic  acid  falls  short  of  100  per  cent 
by  a  small  fraction  of  i  per  cent. 

426.  A  Third  Example  of  Class  II:  Action  of  NaOH  on 
NH4C1. — Another  important  example  of  Class  II  is  found  in 
the  action  of  sodium  hydroxide  and  ammonium  chloride.  The 
addition  of  dilute  NaOH  to  a  solution  of  NH4C1  does  not  produce 
any  visible  effect;  but  evidence  that  the  reaction 

NaOH+NH4Cl±;NaCl+NH4OH 

takes  place  may  be  obtained  in  two  ways:  first,  by  finding  a 
great  decrease  in  conductivity  on  mixing  superimposed  layers 


Applications  of  the  Ionic  Hypothesis  261 

of  the  two  solutions;  and  secondly,  by  noting  the  odor  of 
ammonia  given  off  by  reason  of  the  partial  dissociation  of  the 
NH4OH  present  in  the  solution 

NH4OH^NH3+H2O. 

The  compound  equation  of  the  reaction  follows: 
NaOH^Na++OH- 


it         It 
NaCl  NH4OH. 

Since  sodium  hydroxide  and  ammonium  chloride  are  highly 
ionized,  and  ammonium  hydroxide  is  little  ionized,  this  reaction 
is  completely  analogous  to  that  between  HC1  and  NaC2H302  : 

HCl+NaC2H3O2^NaCl+HC2H302. 

Each  reaction  takes  place  nearly  completely  from  left  to  right 
because  one  product  is  but  little  ionized.  The  graphs  for 
this  reaction,  Figs.  62  and  63,  are  closely  similar  to  those  for 


~  -  NH.CI 


FIG.  62  FIG.  63 

neutralization,  Figs.  58  and  59,  and  for  the  reaction  between 
HC1  and  NaC2H302,  Figs.  60  and  61. 

427.  Summary  of  Class  II  Reactions. — As  we  have  pointed 
out,  all  these  reactions  are  alike,  in  that  two  highly  ionized 
electrolytes  react  to  form  one  highly  ionized  electrolyte  and  one 
little  ionized  electrolyte.  Invariably  reactions  of  this  class  are 
nearly  complete.  The  smaller  the  degree  of  ionization  of  the 
little  ionized  product,  the  more  completely  the  reaction  takes. 


262  Introduction  to  General  Chemistry 

In  the  resulting  mixture  the  little  ionized  substance  is  present, 
of  course  almost  wholly  in  the  molecular  form. 

428.  Strength  of  Acids. — Since  all  salts  are  highly  ionized, 
the  reaction  between  any  highly  ionized  acid  and  a  salt  of  a 
little  ionized  acid  must  belong  to  Class  II.     We  may  therefore 
predict  that,  as  in  the  second  example  studied,  such  reactions 
will  give  nearly  100  per  cent  yields  of  their  products,  and  that 
in  the  resulting  solution  there  will  be  present  the  little  ionized 
acid  instead  of  the  highly  ionized  acid  originally  used.     The 
highly  ionized  acid  may  be  said  to  have  displaced  the  little 
ionized  acid  from  its  salt.     As  a  result,  we  may  call  the  former 
a  strong  acid  and  the  latter  a  weak  acid,  and  may  say  that  a 
strong  acid  always  displaces  a  weak  acid  from  its  salts. 

429.  Strength  of  Bases. — Just  as  we  call  a  highly  ionized 
acid  a  strong  acid  and  a  little  ionized  acid  a  weak  acid,  so  we 
may  call  a  highly  ionized  base  a  strong  base  and  a  little  ionized 
base  a  weak  base.     Since  all  reactions  between  strong  bases  and 
the  salts  of  weak  bases  (see  third  example,  426)  are  examples  of 
Class  II,  we  can  predict  that  the  yield  of  weak  base  and  salt  of 
the  strong  acid  will  be  nearly  100  per  cent.     In  other  words,  a 
strong  base  will  always  displace  a  weak  base  from  its  salt. 

430.  Two  Useful  Laws. — The  foregoing  law  and  that  given 
in  the  paragraph  on  the  strength  of  acids  (428)  have  been  of 
very  great  practical  convenience  to  chemists.     These  laws  fail 
only  when  the  salt  of  the  weak  acid  is  little  ionized,  a  case  so 
rare  that  the  usefulness  of  the  rules  is  virtually  unaffected.    The 
laws  are  of  course  only  special  cases  of  the  fundamental  one  that 
if  two  highly  ionized  substances  react  to  form  one  little  ionized  sub- 
stance and  one  highly  ionized  substance,  the  reaction  will  be  nearly 
complete. 

431.  Suppression  of  the  lonization  of  a  Weak  Acid  or  a 
Weak  Base. — Since  the  strength  of  an  acid  or  a  base  is  deter- 
mined by  its  tendency  to  ionize,  any  factor  that  has  an  influence 
on  this  tendency  will  affect  the  strength  or  weakness  of  the  acid 
or  base.    We  must  now  consider  this  important  subject  and  will 
begin  by  studying  the  action  of  NH4C1  on  a  solution  of  the  weak 
base  NH4OH. 


Applications  of  the  Ionic  Hypothesis  263 

If  we  add  a  little  phenolphthalein  to  very  dilute  NH4OH  a 
bright,  red-colored  solution  results.  This  shows  that  the  solu- 
tion is  alkaline,  and  therefore  that  it  contains  an  abundance  of 
OH~  ions.  Upon  addition  of  a  little  NH4C1  to  this  red  solution 
the  color  disappears  almost  completely.  This  proves  that  the 
number  of  OH~  ions  present  has  been  very  greatly  decreased.  In 
order  to  understand  how  this  has  happened,  we  must  consider 
the  matter  from  the  standpoint  of  ionic  equilibrium.  A  solu- 
tion of  NH4OH  is  ionized  to  a  small  extent,  thus  : 

NH4OH^NH4++OH-. 

Ammonium  chloride,  on  the  other  hand,  is  very  highly  ionized: 
NH4CteNH4++Cl-. 

If  then  we  add  an  equivalent  amount  of  NH4C1  to  a  dilute  solu- 
tion of  NH4OH,  the  number  of  NH4+  ions  per  cubic  centimeter 
will  be  increased  many  fold.  The  OH~  ions  present  will  therefore 
collide  with  NH4+  ions  and  combine  with  them  far  more  fre- 
quently than  before.  Since  the  rate  of  dissociation  of  NH4OH 
molecules  into  ions  is  not  affected  by  the  presence  of  the  NH4C1, 
this  increased  rate  of  union  of  NH4+  and  OH"  ions  causes  a 
great  shift  to  the  left  of  the  equilibrium 


For  example,  it  has  been  found,  by  methods  that  we  need  not 
consider  here,  that  the  addition  of  i  g.  of  NH4C1  to  100  c.c.  of 


FIG.  64  FIG.  65 

decinormal  NH4OH  will  decrease  the  number  of  OH"  ions 
present  about  one  hundred  fold.  In  other  words,  the  ionization 
of  the  base  will  be  decreased  one  hundred  fold  (see  Figs.  64  and 
65).  We  may  now  state  the  general  law  of  which  the  case  just 


264  Introduction  to  General  Chemistry 

studied  is  a  typical  example:  The  ionization  of  a  weak  base  is 
greatly  suppressed  by  the  addition  of  a  salt  of  the  base.  This 
means  that  a  weak  base  is  made  still  weaker  by  the  addition  of  its 
soluble  salts. 

In  a  similar  manner  the  ionization  of  a  weak  acid  is  greatly 
suppressed  by  the  addition  of  any  of  its  soluble  salts;  that  is,  a  weak 
acid  is  made  still  weaker  by  adding  one  of  its  salts.  For  example, 
the  addition  of  NaC2H3O2  to  a  red  solution  of  acetic  acid, 
HC2H302,  containing  litmus  changes  the  color  from  red  to 
purple,  thus  showing  a  great  decrease  in  the  number  of  H+  ions, 
and  therefore  a  great  decrease  in  ionization  of  the  acid. 

432.  The  Common  Ion  Law.  —  A  base  and  any  one  of  its 
salts  must  of  necessity  have  one  ion  in  common.     (The  NH4+ 
ion  is  common  to  NH4OH  and  NH4C1.)    An  acid  also  must  have 
one  ion  in  common  with  any  of  its  salts.    We  may  therefore 
state  the  principle  of  the  foregoing  laws  as  follows:  Suppression 
of  the  ionization  of  a  little  ionized  substance  occurs  when  we  add  to 
its  solution  a  highly  ionized  substance  having  a  common  ion.    This 
is  the  Common  Ion  Law,  a  very  important  generalization.    The 
examples  already  cited  are  by  no  means  the  only  ones  of  impor- 
tance.   For  example,  it  is  plain  that  the  ionization  of  NH4OH 
must  be  suppressed  by  the  addition  of  NaOH  or  KOH  because 
of  the  increase  in  concentration  of  the  common  OH~  ion;   and 
that  the  ionization  of  HC2H302  must  likewise  be  suppressed  by 
the  addition  of  any  strong  acid  like  HC1  or  HN03.    The  effect 
of  a  highly  ionized  substance  on  the  ionization  of  another  highly 
ionized  substance  having  one  ion  in  common  is  of  the  same  type 
but  very  much  smaller  in  degree  than  when  the  second  substance 
is  slightly  ionized. 

We  shall  next  consider  the  application  of  the  Common  Ion 
La.w  to  solutions  of  acids  and  bases  and  thus  obtain  a  definition 
of  the  term  neutrality. 

433.  Neutrality  Defined.  —  We  have  already  learned  (420) 
that  water  is  slightly  ionized,  thus, 


Each  cubic  centimeter  of  pure  water  must  therefore  contain 
exactly  as  many  H+  as  OH~  ions.     Since  all  acids  give  H+ 


Applications  of  the  Ionic  Hypothesis  265 

ions,  the  addition  of  an  acid  to  water,  in  accord  with  the  common 
ion  law,  will  greatly  suppress  the  ionization  of  water.  Therefore 
acid  solutions  will  contain  far  less  OH~  ions  per  cubic  centimeter 
than  pure  water.  In  an  acid  solution  the  number  of  H+  ions 
greatly  exceeds  the  number  of  OH~  ions.  The  ionization  of 
water  is  also  greatly  suppressed  by  the  addition  of  a  base,  since 
all  bases  have  OH~  ions  in  common  with  water.  In  basic  solu- 
tions the  number  of  H+  ions  per  cubic  centimeter  is  far  less  than 
in  pure  water  and  therefore  the  number  of  OH~  ions  greatly 
exceeds  the  number  of  H+  ions.  Since  we  may  consider  water 
a  typically  neutral  substance  we  may  define  a  neutral  solution  as 
one  in  which  the  number  of  H+  ions  equals  the  number  of  OH~ 
ions.  Since,  as  we  have  already  learned,  a  strong  acid  completely 
neutralizes  a  strong  base,  as  for  example  in  the  reaction 

HCl+NaOH±?NaCl+HaO, 

we  conclude  that  in  the  resulting  solution  the  number  of  H+  ions 
is  just  equal  to  the  number  of  OH~  ions:  this  is  the  criterion 
of  complete  neutrality. 

434.  First  Example  of  Class  III :  The  Action  of  a  Weak  Acid 
on  a  Strong  Base. — Under  Class  III  we  shall  study  reactions  in 
which  one  little  ionized  and  one  highly  ionized  substance  give 
products,  one  of  which  is  little  ionized,  the  other  highly  ionized. 
As  the  first  example  we  shall  study  the  reaction  between  little 
ionized  acetic  acid  (a  weak  acid)  and  highly  ionized  sodium 
hydroxide  (a  strong  base).  These  react  thus: 

HC2H3O2+ NaOH  ±>  H2O+ NaC2H3O2. 

Of  the  products,  water  is  very  slightly  ionized,  while  sodium 
acetate,  ,NaC2H3O2,  is  highly  ionized.  If  we  mix  equal  volumes 
of  normal  solutions  of  HC2H3O2  and  NaOH,  that  is,  if  we  add 
to  the  NaOH  solution  exactly  that  quantity  of  acetic  acid  that 
would  neutralize  it  if  the  reaction  were  complete,  we  find  that 
the  resulting  mixture  is  not  neutral  but  is  still  alkaline  to  litmus. 
The  fact  that  the  mixture  is  alkaline  means  that  the  number  of 
H0~  ions  is  greater  than  the  number  of  H+  ions  present.  The 


266  Introduction  to  General  Chemistry 

cause  of  this  condition  is  most  easily  understood  by  aid  of  the 
compound  equation 

HC2H302^H++C2H302- 


It         it 
H2O    NaC2H3O2 

and  Figs.  66  and  67.  At  the  instant  of  mixing,  the  solution 
contains  an  abundance  of  OH~  ions  (Fig.  66);  these  reduce 
greatly  the  number  of  H"1"  ions  present  by  forming  H20  molecules  : 


The  removal  of  H+  ions  disturbs  the  equilibrium 
HC2H302±>H++C2H302-, 

which  shifts  greatly  to  the  right,  thus  producing  both  H+  and 
C2H3O2~  ions.     While  the  former  unite  with  OH~  almost  (but 


HC,H,0, 


FIG.  66  FIG.  67 

not  quite)  completely,  the  latter  remain  for  the  larger  part  free 
in  the  solution,  and  by  their  great  tendency  to  unite  again  with 
H+  ions  to  form  little  ionized  HC2H302  serve  still  further  to 
diminish  the  number  of  free  H+  ions.  In  the  final  equilibrium 
mixture,  shown  in  Fig.  67,  the  number  of  OH"  ions  is  greater 
than  the  number  of  H+  ions  because  of  the  great  tendency  of  the 
latter  to  unite  readily  with  either  C2H3O2~  ions  or  OH"  ions. 
That  the  OH"  ions  get  by  far  the  lion's  share  of  the  H+  ions  is 
owing  to  the  fact  that  water  is  much  less  ionized  than  acetic  acid. 


Applications  to  the  Ionic  Hypothesis  267 

Since  the  mixture  contains  more  OH~  than  H+  ions  (see  Fig. 
67)  it  is  not  neutral  but  alkaline. 

435.  The  Action  of  Water  on  Sodium  Acetate. — In  the  fore- 
going paragraph  we  studied  the  equilibrium 

HC2H3O2+NaOH±5NaC2H302+H2O. 

The  composition  of  the  equilibrium  solution  was  shown  in  Fig.  67. 
It  must  be  plain  from  the  deduction  of  281,  that  exactly  the  same 
equilibrium  solution  would  be  obtained  if  we  should  dissolve  in 
the  same  quantity  of  water  pure  sodium  acetate,  NaC2H302,  in 
exactly  the  amount  that  would  be  produced  by  the  complete 
union  of  all  the  HC2H3O2  and  NaOH  used  in  the  first  case.  As  a 
matter  of  fact  we  find  that  a  solution  of  pure  sodium  acetate  is  not 
neutral  but  alkaline  to  litmus.  The  action  of  H2O  on  NaC2H302 
takes  place  thus:  the  salt  first  dissolves  and  at  once  ionizes 
highly  to  form  many  Na+  and  C2H3O2~  ions.  Water,  although 
but  slightly  ionized,  contains  some  H+  and  OH"  ions.  Occasional 
collisions  of  H+  and  C2H3O2~  ions  will  occur,  and  part  of  these 
collisions  will  result  in  unions  to  form  HC2H302  molecules;  and 
as  the  latter  have  but  little  tendency  to  ionize,  the  result  is  a 
great  decrease  in  the  number  of  H+  ions  present.  This  in  turn 
disturbs  the  equilibrium 

H2O±>H++OH-, 

which  in  consequence  shifts  to  the  right  and  so  brings  more  OH~ 
ions  into  the  solution.  A  few  but  not  many  of  the  OH"  ions 
unite  with  Na+  ions  to  form  molecules  of  NaOH,  but  most  of 
the  OH"  ions  remain  free,  thus  producing  in  the  solution  a 
decided  excess  of  OH"  ions  over  H~*~  ions  (see  Fig.  67),  and  so 
making  the  solution  alkaline  to  litmus.  Briefly  stated,  water 
acts  on  sodium  acetate  to  a  small  extent,  thus, 

NaC2H302+H20±>HC2H302+NaOH, 

t 

and  since  NaOH  is  highly  ionized,  while  HC2H3O2  is  little  ionized, 
the  reaction  of  the  solution  is  alkaline.  The  composition  of  a 
water  solution  of  sodium  acetate  is  that  shown  in  Fig.  67. 

436.  Hydrolysis  of  Salts. — The  soluble  salts  of  all  weak  acids 
with  the  strong  bases  sodium,  potassium,  calcium,  or  barium 


268  Introduction  to  General  Chemistry 

hydroxide  give  alkaline  solutions  when  dissolved  in  water.  In 
every  case  the  reason  is  the  same  as  that  given  for  the  alkaline 
reaction  of  sodium  acetate  solution.  The  effect  of  water  on  the 
salt  of  a  weak  acid  and  a  strong  base  is  an  example  of  the  type 
of  reaction  called  hydrolysis  (or  also  hydrolytic  dissociation). 
Hydrolysis  may  be  defined  as  a  double  decomposition  reaction 
in  which  water  is  one  of  the  reacting  substances.  The  solutions 
of  salts  of  all  weak  acids  and  strong  bases  are  alkaline  in  reaction. 
Other  things  being  equal,  the  weaker  the  acid  from  which  the  salt 
is  derived  the  greater  the  extent  of  the  hydrolysis;  that  is,  the 
greater  the  alkalinity  of  the  solution. 

On  the  other  hand,  some  salts  (other  than  acid  salts  like 
NaHS04)  give  solutions  that  have  an  acid  reaction  (176). 
Among  such  are  the  chlorides,  sulfates,  and  nitrates  of  copper, 
lead,  iron,  zinc,  aluminum,  etc.  Experiments  show  that  the 
hydroxides  of  all  these  elements  are  weak  bases.  It  would 
therefore  seem  probable  that  the  acidity  ot  solutions  of  the  salts 
of  these  bases  with  strong  acids  is  due  to  hydrolysis,  and  that 
the  behavior  of  such  salts  with  water  is  the  counterpart  of  the 
behavior  of  salts  of  weak  acids  with  strong  bases. 

437.  A  Second  Example  of  Class  III  :  The  Action  of  a  Strong 
Acid  on  a  Weak  Base.  —  The  action  of  a  strong  acid  on  a  weak 
base  is  plainly  the  reverse  of  that  just  discussed:  the  action  of 
water  on  the  salt  of  a  strong  acid  and  weak  base.  It  follows  that 
a  weak  base  does  not  react  completely  with  the  theoretical  or 
chemically  equivalent  amount  of  a  strong  acid,  and  in  con- 
sequence the  resulting  mixture  is  still  acid  in  its  reaction.  The 
action  of  HC1  on  the  weak  base  NH4OH  will  serve  as  a  simple 

illustration: 

NH4OH^OH-+NH4+ 


it        It 
H2O    NH4CL 

Comparison  of  this  reaction  with  that  for  HC2H302  and  NaOH 
where  we  have  a  weak  acid  and  strong  base  will  bring  out  com- 
plete analogy.  Experiment  shows  that  a  solution  of  NH4C1  in 
water  is  not  neutral  but  slightly  acid  in  reaction.  Briefly 


Applications  of  the  Ionic  Hypothesis  269 

stated,  NH4OH  does  not  completely  neutralize  an  equivalent 
amount  of  HC1  because  it  is  a  weak  base.  Conversely,  water 
acts  on  pure  NH4C1  to  form  some  free  HC1  and  NH4OH. 

438.  Class  IV:  The  Action  of  a  Weak  Acid  and  a  Weak 
Base. — Under  Class  IV  we  shall  include  reactions  between  two 
little  ionised  substances,  which  give  as  products  one  little  ionized 
and  one  highly  ionized  substance.  The  only  reactions  of  im- 
portance in  Class  IV  are  those  between  a  weak  acid  and  a  weak 
base,  the  products  being  water  and  a  salt.  Acetic  acid  and 
ammonium  hydroxide  are  both  moderately  weak  (but  not 
extremely  weak) .  They  react  thus : 

HC2H3O2+NH4OH^NH4C2H302+H20. 

The  reaction  is  not  complete,  as  in  the  case  of  the  action  of  a 
strong  acid  and  a  strong  base,  but  reaches  equilibrium  when  a 
few  tenths  of  i  per  cent  of  the  free  un-ionized  acid  and  free 
un-ionized  base  are  still  present  in  the  solution.  The  conditions 
before  and  after  the  reactions  are  shown  in  Figs.  68  and  69. 


\HC,H,0, 


HC,H,0, 


FIG.  68  FIG.  69 

If  on  the  other  hand  the  solution  is  made  by  dissolving  solid 
NH4C2H3O2  in  water  partial  hydrolysis  takes  place,  giving  a 
mixture  the  composition  of  which  is  also  represented  by  Fig.  69. 

If  both  acid  and  base  are  extremely  weak  the  extent  of  the 
hydrolysis  will  be  much  greater  than  in  the  case  of  NH4C2H3O2. 
In  fact,  in  such  cases  hydrolysis  may  be  so  nearly  complete  that 
we  may  say  that  extremely  weak  acids  in  water  solution  do  not 
form  salts  with  extremely  weak  bases. 


270  Introduction  to  General  Chemistry 

439.  Heat  of  lonization. — The  heat  liberated  or  absorbed  by 
the  complete  dissociation  into  its  ions  of  one  formula  weight 
of  a  dissolved  electrolyte  is  called  its  heat  of  ionization.  In 
some  cases  heat  is  absorbed,  in  other  cases  it  is  liberated,  when 
the  substances  are  ionized,  but  in  the  great  majority  of  cases 
the  heat  of  ionization  is  very  small.  For  practical  purposes  we 
may  say  that  the  heat  of  ionization  of  readily  ionizable  elec- 
trolytes is  almost  negligible.  Little  ionized  substances  often 
have  appreciable  heats  of  ionization.  This  is  notably  the  case 
with  water,  for  which  we  have  the  following, 

H++OH-->H2O+ 13,700  cal. 

It  was  stated  earlier  (362)  that  the  heat  of  neutralization  of 
a  strong  acid  by  a  strong  base  is  almost  the  same  in  all  cases, 
namely  13,700  calories.  The  reason  can  now  be  seen.  We 
know  that  in  the  neutralization  of  a  strong  acid  by  a  strong  base 
in  dilute  solution  the  principal  change  is  the  union  of  H+  and  OH~ 
ions  to  form  water.  In  other  words,  the  simplified  equation  of 

neutralization  is 

H++OH-->H20. 

Since  strong  acids  and  bases,  as  well  as  most  salts,  have 
negligible  heats  of  ionization;  and  since,  moreover,  very  little 
dissociation  or  union  of  ions,  other  than  H+  and  OH~,  occurs 
in  neutralization  (421,  Fig.  59),  the  heat  produced  in  the  reaction 
is  simply  that  due  to  the  formation  of  water  from  its  ions.  It 
is  for  this  reason  that  heats  of  neutralization  are  practically 
the  same  for  all  strong  acids  and  bases:  13,700  cal.  for  one 
formula  weight  (18  g.)  of  water  formed. 

The  heat  of  neutralization  of  ammonium  hydroxide  by  a 
strong  acid  is  12,300  calories.  The  difference,  13,700—12,300  = 
1,400  cal.,  is  the  heat  of  ionization  of  the  weak  base. 

In  reactions  between  solutions  of  two  highly  ionized  salts 
which  form  by  interaction  two  other  highly  ionized  and  easily 
soluble  salts  no  appreciable  heat  change  is  observed.  This  is 
because  in  such  reactions  very  little  change  takes  place  (418, 
Fig.  55),  and  such  changes  as  do  occur  are  accompanied  by  nearly 
negligible  heats  of  ionization. 


Applications  of  the  Ionic  Hypothesis  271 

440.  Indicators. — In  addition  to  litmus,  which  is  used  so 
often  to  indicate  the  acidity  or  alkalinity  of  solutions,  a  number  of 
other  colored  substances  are  also  employed.  These  are  called 
indicators.  The  more  important  indicators  besides  litmus  are 
phenolphthalein  and  methyl  orange.  The  former  is  a  colorless 
substance  which  gives  a  bright  red  solution  with  alkalies. 
Methyl  orange  is  orange  color  in  neutral  solution,  pink  in  acid, 
and  yellow  in  alkaline  solution.  In  general,  indicators  are  very 
complex  chemical  substances  whose  formulae  need  not  be  con- 
sidered at  present. 

Since  acid  solutions  always  contain  H+  ions  and  alkaline 
solutions  OH~t  ions,  we  may  say  that  an  indicator  is  a  substance 
which  has  one  color  in  the  presence  of  an  excess  of  H+  ions  and 
a  different  color  in  the  presence  of  an  excess  of  0  H~  ions.  We 
might  expect  that  every  indicator  would  show  its  transition 
shade  of  color  in  an  exactly  neutral  solution;  that  is,  in  a  solu- 
tion where  the  number  of  H+  ions  equals  the  number  of  OH" 
ions.  This,  however,  is  not  the  case.  In  other  words,  most 
indicators  do  not  indicate  perfect  neutrality.  Litmus  is  a  nearly 
perfect  indicator,  but  phenolphthalein  shows  a  change  of  color 
when  the  number  of  OH~  ions  equals  eighty  times  the  number ' 
of  H+  ions;  that  is,  if  a  solution  contains  more  than  eighty 
times  as  many  OH~  as  H+  ions  it  colors  phenolphthalein  red 
(the  alkaline  color) ;  if  it  contains  less  than  eighty  times  as  many 
OH""  ions  as  H+  ions  it  leaves  phenolphthalein  colorless.  On 
the  other  hand,  methyl  orange  shows  an  orange  color  (its  inter- 
mediate shade  between  pink,  the  acid  color,  and  yellow,  the 
alkaline  color)  when  the  number  of  H+  ions  is  about  a  million 
tunes  the  number  of  OH~  ions.  Anomalous  as  it  may  seem  at 
first  thought,  it  is  really  fortunate  that  many  of  our  indicators 
do  not  indicate  perfect  neutrality;  for  suppose  we  wish  to  dis- 
cover how  much  acetic  acid  a  certain  solution  contains.  We 
may  titrate  it  accurately  with  normal  or  decinormal  sodium 
hydroxide  or  other  strong  base  if  we  use  the  right  indicator 
(137).  Now  we  have  learned  that  when  acetic  acid  is  mixed  with 
exactly  the  theoretically  equivalent  amount  of  NaOH  the  result- 
ing solution  is  not  perfectly  neutral  but  in  reality  slightly  alkaline 


272 


Introduction  to  General  Chemistry 


(434) .  In  accord  with  this  we  found  that  a  solution  of  NaC2H3O2 
was  slightly  alkaline  to  litmus,  showing  that  the  number  of  OH~ 
ions  was  greater  than  the  number  of  H+  ions.  Therefore  we 
must  use  as  a  titration  indicator  one  which  shows  its  change  of 
color  when  the  number  of  OH~~  is  greater  than  the  number  of  H+ 
ions.  We  find  that  phenolphthalein  proves  to  be  just  right  for 
the  purpose.  In  general,  we  use  phenolphthalein  as  indicator 
in  titrating  all  moderately  weak  acids. 

If  we  wish  to  titrate  NH4OH  with  HC1  we  cannot  use 
phenolphthalein,  because  a  solution  of  NH4C1  contains  more  H+ 
than  OH~  ions.  Such  a  solution  "seems"  acid  to  this  indicator. 
We  must  use  one  which  changes  color  when  the  number  of  H+ 
ions  exceeds  the  number  of  OH~  ions,  and  for  this  case  we  find 
methyl  orange  satisfactory.  In  general,  we  use  methyl  orange 
in  titrating  moderately  weak  bases.  The  acid  used  in  such 
titrations  must  be  a  strong  one.  In  titrating  a  strong  base  with 
a  strong  acid  any  of  these  indicators  gives  sufficiently  accurate 
results.  Table  XVIII  gives  the  colors  of  indicators  in  solutions 

TABLE  XVIII 


Hydrogen  ion  concentration  

io-3 

io-5 

10-7 

io-» 

Hydroxyl  ion  concentration  

io-" 

io-9 

10-7 

io-6 

Methyl  orange  

Pink 

Yellow 

Yellow 

Yellow 

Litmus 

Red 

Red 

Purple 

Blue 

Phenolphthalein 

Colorless 

Colorless 

Colorless 

Red 

of  hydrogen  and  hydroxyl  ion  concentrations  near  those  at 
which  the  color  change  occurs.  In  this  table  the  concentrations 
are  given  in  gram  molecular  weights  per  liter.  If  the  H+  con- 
centration is  io~3,  1,000  liters  contain  i  g.  of  H+  ion. 

441.  Summary  on  Equilibrium  between  Soluble  Electro- 
lytes.— If  we  mix  solutions  of  two  electrolytes,  AB  and  CD, 
having  no  ion  in  common,  a  double  decomposition  reaction, 

AB+CD^AD+CB, 

takes  place  to  a  greater  or  less  extent,  because  of  the  tendency 
of  each  positive  ion  to  combine  with  each  negative  ion  present. 


Applications  of  the  Ionic  Hypothesis  273 

If  all  four  substances  of  the  preceding  equation  are  highly 
ionized  (Class  I,  415),  the  mixed  solution  will  contain  largely 
the  four  sorts  of  free  ions,  A+,  B~7  C+,  andZ>~.  Only  a  small 
percentage  of  the  dissolved  material  will  be  present  as  molecules. 
Figure  56  represents  equilibrium  conditions  in  an  example  in 
which  equivalent  amounts  of  initial  materials  were  used. 

If  one  of  the  four  substances  (say  AD)  is  little  ionized  (Class 
II,  419),  then  the  large  numbers  of  A+  and  D~~  ions  shown  in 
Fig.  56  cannot  exist  side  by  side  in  the  mixed  solution,  since  they 
will  very  largely  combine  to  form  AD  molecules.  The  disap- 
pearance of  A+  and  D~  ions  allows  AB  and  CD  molecules  more 
or  less  completely  to  dissociate.  The  final  result,  shown  in 
Figs.  59,  61,  and  63,  is  a  nearly  complete  reaction,  AD  being 
present  almost  wholly  in  un- ionized  form  and  CD  to  a  small 
extent  as  molecules,  but  largely  as  C+  and  D~  ions. 

A  generalization  of  much  importance  is  found  in  the  Common 
Ion  Law:  suppression  of  the  ionization  of  a  little  ionize  sub- 
stance occurs  when  we  add  to  its  solution  a  highly  ionized 
substance  having  a  common  ion. 

Since  in  pure  water  the  number  of  H+  ions  is  equal  to  the 
number  of  OH~  ions,  and  since  we  may  consider  pure  water  a 
perfectly  neutral  substance,  we  define  a  neutral  solution  as  one 
in  which  the  number  of  H+  ions  is  exactly  equal  to  the  number 
of  OH~  ions. 

Class  III  (434,  437)  comprises  reactions  in  which  one  little 
ionized  substance  reacts  with  a  highly  ionized  substance  to 
form  products  one  of  which  is  slightly,  the  other  highly,  ionized. 
Examples  are  found  in  the  neutralization  of  a  weak  acid  by  a 
strong  base;  or  of  a  weak  base  by  a  strong  acid.  In  such  cases 
the  reaction  is  more  or  less  incomplete.  The  weaker  the  acid 
or  base  taken,  the  less  complete  is  the  neutralization.  Con- 
versely, salts  of  weak  acids  or  of  weak  bases  are  hydrolyzed  by 
water.  The  former  give  solutions  which  are  alkaline,  the  latter 
those  which  are  acid,  in  reaction.  This  kind  of  action  is  called 
hydrolytic  dissociation. 

Under  Class  IV  (438)  it  was  pointed  out  that  weak  acids  and 
weak  bases  always  react  incompletely,  and  that  when  either 


274  Introduction  to  General  Chemistry 

or  both  are  extremely  weak,  salt  formation  may  not  occur  in 
solution  (177).     . 

We  have  seen  that  indicators  change  color  according  to  the 
concentration  of  H+  and  OH~  ions  present.  Litmus  shows 
its  neutral  tint  when  the  numbers  of  H+  and  OH~  ions  are 
nearly  equal.  Phenolphthalein  requires  an  excess  of  OH~ 
ions  to  change  color,  while  methyl  orange  requires  an  excess  of 
H+  ions. 


CHAPTER  XIX 

APPLICATIONS  OF  THE  IONIC  HYPOTHESIS.    REACTIONS 
INVOLVING  CHANGES  OF  STATE 

442.  Introduction.  —  In  the  present  chapter  we  shall  study 
precipitation  from  the  standpoint  of  the  ionic  hypothesis  in 
order  to  understand  the  underlying  principles  of  this  most 
important  means  of  separating  substances.  In  equations  for 
precipitation  reactions,  the  substance  precipitated  will  be  indi- 
cated by  a  downward-pointing  arrow. 

If  we  consider  the  familiar  examples  of  precipitation  repre- 
sented by  the  following  equations, 


we  might  conclude  that  AgCl  and  BaSO4  are  precipitated  because 
they  are  insoluble  in  water.  We  might  even  be  tempted  to  say 
that  in  the  reaction 

AB+CD^AD+CB, 

if  either  AD  or  CB  is  an  insoluble  substance  it  will  be  pre- 
cipitated. This  statement  contains  something  of  the  truth, 
but  it  is  far  from  the  whole  truth,  as  the  following  examples  will 
prove.  Calcium  carbonate,  CaC03  (marble),  is  an  almost 
insoluble  substance.  If  we  mix  solutions  of  calcium  chloride 
and  carbonic  acid  we  might  expect  to  get  a  precipitate  of  calcium 
carbonate,  thus, 

CaCl2+H2CO3^CaC03i+2HCL 

Not  a  trace  of  precipitate  is  formed.  On  the  other  hand  potas- 
sium chlorate,  KC103,  is  easily  soluble  in  water;  but  if  we  add 
a  saturated  solution  of  potassium  bromide,  KBr,  to  a  saturated 
solution  of  sodium  chlorate,  NaClO3,  a  precipitate  of  KC103 
forms.  Evidently  the  matter  is  not  as  simple  as  at  first  thought 
it  appears  to  be.  The  separation  of  a  solid  from  a  solution  is 

275 


276  Introduction  to  General  Chemistry 

obviously  the  reverse  of  the  passage  of  a  solid  into  solution. 
Accordingly,  in  beginning  the  study  of  precipitation,  it  will  be 
advisable  for  the  student  to  read  again  sections  120-23.  In 
section  122  it  is  stated,  "A  solution  which  at  a  fixed  tempera- 
ture will  dissolve  no  more  of  a  given  substance  is  called  a 
saturated  solution.  When  we  speak  of  the  solubility  of  a  sub- 
stance we  mean  the  amount  of  substance  dissolved  in  a  given 
amount  of  water  in  the  case  of  the  saturated  solution." 

443.  The  Kinetic  Theory  of  Solution.  —  When  a  soluble  salt 
is  brought  into  water  its  molecules  begin  to  leave  the  surface 
of  the  solid  and  pass  into  the  water.  Immediately  thereafter 
dissolved  salt  molecules  will  occasionally  strike  the  surface 
of  the  solid  and  in  some  cases  remain  attached  thereto.  Finally, 
when  the  solution  has  become  saturated  we  may  imagine  that 
the  equilibrium  between  dissolved  and  solid  salt  is  the  result 
of  the  passage  of  molecules  into  and  out  of  solution  at  exactly 

equal  rates,  thus: 

AB^AB 
Solid    Dissolved 

This  picture  is,  however,  incomplete,  since  the  salt  is  partly 
ionized.  The  dissolved  molecules  are  therefore  in  equilibrium 
with  their  ions  as  well  as  with  the  solid  salt,  thus: 


Solid        Dissolved 

444.  Graphic    Representation    of    a    Solid    Electrolyte    in 
Equilibrium  with  Its  Saturated  Solution.  —  We  shall  represent 


NaCl 


FIG.  70 

a  solid  electrolyte  (acid,  base,  or  salt)  graphically  by  a  cross- 
hatched  square.  The  condition  of  a  saturated  solution  of  a 
soluble  salt  (NaCl,  for  example)  in  contact  with  an  excess  of 
the  solid  salt  may  then  be  represented  as  in  Fig.  70. 


Applications  of  the  Ionic  Hypothesis  277 

445.  The  Solubility  of  Molecules.     Molecular  Solubility.— 

If  we  except  a  small  number  of  electrolytes  like  sulfuric  and 
nitric  acids,  which  mix  with  water  in  all  proportions,  all  other 
acids,  bases,  and  salts  have  limited  solubilities  in  water.  Since 
all  electrolytes  are  more  or  less  ionized  in  solution,  the  dissolved 
substance  is  present  partly  as  molecules  and  partly  as  ions. 
Therefore  the  total  solubility  of  a  substance  in  a  solution  saturated 
at  a  given  temperature  is  the  sum  of  the  solubility  of  its  molecules 
and  the  solubility  of  its  ions.  It  seems  reasonable  to  assume 
that  the  limited  solubility  of  an  electrolyte  as  a  whole  is  the 
result  of  the  limited  solubility  of  its  molecules  rather  than  of  its 
ions.  Two  reasons  may  be  given  for  this  assumption  which  will 
be  amply  confirmed  by  additional  evidence  to  be  considered 
later. 

In  the  first  place  the  solid  salt  passes  into  and  out  of  solution 
as  molecules  (see  Fig.  70).  If  the  molecules  have  a  limited 
solubility,  this  would  limit  the  solubility  of  the  ions  as  well,  since 
the  latter  and  the  former  are  directly  in  equilibrium.  Therefore 
it  is  sufficient  to  assume  limited  solubility  of  the  molecules  in  order 
to  explain  limited  total  solubility.  Secondly,  the  small  solubility 
of  a  difficultly  soluble  salt  like  CaS04  (100  c.c.  of  water  dissolve 
o.25g.  of  CaS04)  cannot  be  due  to  a  correspondingly  small 
solubility  of  Ca++  or  S04  ions,  since  solutions  of  CaCl2  and 
many  other  easily  soluble  and  highly  ionized  calcium  salts  con- 
tain an  abundance  of  Ca++  ions,  and  solutions  of  H2S04  and 
many  easily  soluble  and  highly  ionized  sulfates  contain  large 
concentrations  of  S04  ions.  We  shall  assume,  therefore,  that 
at  a  given  temperature  the  solubility  of  an  acid,  base,  or  salt  is 
limited  by  the  solubility  of  its  molecules;  and  we  shall  call  the 
solubility  of  the  molecules  (in  the  saturated  solution)  the 
molecular  solubility  (abbreviated  M.S.)  of  the  substance.  Sum- 
marizing, we  may  say  that  when  a  solid  electrolyte  is  mixed  with 
water  at  a  fixed  temperature  the  substance  dissolves  and  the 
concentration  of  the  solution  increases  until  the  M.S.  is  reached; 
the  solution  is  then  saturated  (at  that  temperature),  and  the 
molecules  are  in  equilibrium  with  the  ions  and  with  the  solid 
substance. 


278  Introduction  to  General  Chemistry 

446.  The  Cause  of  Precipitation. — We  are  now  ready  to 
apply  the  foregoing  principles,  together  with  those  learned  in 
chapter  xviii,  to  the  process  of  precipitation.     We  have  learned 
(414)  that  in  the  reaction 

AB+CD±;AD+CB, 

if  all  four  substances  are  easily  soluble  and  highly  ionized  the 
resulting  solution  contains  largely  the  four  sorts  of  ions,  together 

_  with  small  proportions  of  the 

four  kinds  of  molecules.    The 
conditions  before  and  after 
the  reaction  are  shown  in  Figs. 
-<flli11^         "XJ)~    55  and  56.    Now  let  us  suppose 
that  one  of  the  products  AD 
is  not  very  soluble,  so  that  its 
molecular  solubility  (M.S.)  is 
•v\CB    ,o   ^^w    less  than  that  corresponding  to 
,'     ^"      /      the  concentration  of  the  mole- 
cules of  AD  formed  in  reaction 
(Fig.    56).     In    this    case   the 
~!AD  amount   of  molecules   of    AD 
formed  in  excess  of  the  M.S. 

will  separate  out  of  solution  as 
FIG.  71 

a   precipitate.     Fig.   71  shows 

the  resulting  condition  for  the  case  where  the  M.S.  of  AD  is 
rather  small  but  not  extremely  small.  By  comparison  of  Figs. 
56  and  71  we  see  that  an  appreciable  shift  in  equilibrium  of  the 
dissolved  substances  accompanies  the  partial  precipitation  of  AD. 

447.  The  Precipitation  of  KC1O3. — An  actual  example  con- 
forming perfectly  to  the  conditions  set  forth  in  the  preceding 
paragraph  is  found  in  the  reaction 

KBr+NaClO3^KClO3|+NaBr. 

Of  the  four  salts,  all  are  very  soluble  except  KC1O3,  which  dis- 
solves only  to  the  extent  of  7  g.  in  100  c.c.  of  water  at  18°.  All 
four  salts  are  highly  and  about  equally  ionized  in  solutions  of 
equal  concentration.  The  conditions  of  the  solutions  of  KBr 


Applications  of  the  Ionic  Hypothesis  279 

and  NaC103  before  mixing  are  shown  with  sufficient  accuracy 
by  Fig.  55,  while  Fig.  56  shows  the  condition  which  the  mixed 
solution  would  reach  if  KC1O3  were  also  very  soluble.  It  hap- 
pens, however,  that  the  amount  of  molecular  KC1O3  which 
tends  to  be  formed  exceeds  the  rather  small  M.S.  of  this  sub- 
stance, and  in  consequence  the  excess  above  the  M.S.  separates 
as  a  precipitate.  Precipitation  continues  until  the  amount  of 
molecular  KC1O3  left  in  solution  is  equal  to  the  M.S.  of  this  sub- 
stance. The  mixture  is  then  in  the  condition  of  equilibrium 
shown  in  Fig.  71.  Comparison  of  Figs.  56  and  71  shows  that 
the  removal  (by  precipitation)  of  KC1O3  from  the  solution  causes 
a  marked  shift  in  the  equilibrium.  We  may  trace  the  stages  as 
follows:  Fig.  56  shows  the  condition  that  would  exist  if  no 
precipitation  occurred.  The  removal  of  KC1O3,  results  in  the 
further  union  of  K+  and  C103~  ions  to  form  more  KC103.  The 
resulting  loss  of  K+  and  C103~  ions  promotes  the  further  ioniza- 
tion  of  KBr  and  NaClO3  respectively  and  thus  increases  the 
numbers  of  Br~  and  Na+  ions.  The  latter  ions  unite  in  part 
to  form  additional  molecular  NaBr.  The  final  result  is  the 
change  from  the  condition  of  Fig.  56  to  that  of  Fig.  71.  The 
principles  here  exemplified  apply  to  all  double  decomposition 
precipitations. 

448.  The  Precipitation  of  CaCO3. — Let  us  now  consider  a 
case  in  which  one  of  the  products  is  precipitated  almost  com- 
pletely. The  reaction 

CaCla+NaaCOj±5CaCO34'+  2NaCl, 

hi  which  CaC03  is  the  precipitate,  will  serve  as  a  typical  illustra- 
tion. Although  CaCO3  appears  to  be  insoluble  in  water,  it  is  in 
fact  slightly  soluble,  and  has  therefore  a  definite  but  very  small 
M.S.  The  other  three  salts,  CaCl2,  Na«CO3,  and  NaCl,  are 
easily  soluble  and  highly  ionized,  and  in  consequence  the  reaction 
between  solutions  of  CaCl2  and  Na2C03  tends  to  reach  the  con- 
dition shown  in  Fig.  56  illustrating  a  Class  I  reaction  (414).  In 
this  respect  it  completely  resembles  the  reaction 


280 


Introduction  to  General  Chemistry 


'"""NCaCl, 


CO," 


It  differs  however  from  this  reaction  in  that  the  M.S.  of  CaC03 
is  extremely  small  compared  with  the  M.S.  of  KC103.  In  con- 
sequence the  CaC03  formed  pre- 
cipitates almost  completely,  as 
illustrated  in  Fig.  72.  In  all 
double  decomposition  reactions 
of  the  above-mentioned  types 
(all  involved  substances  highly 
ionized)  the  precipitation  is  the 
more  complete  the  less  the  M.S. 
of  the  precipitate. 

449.  The  Action  of  H2CO3 
on  CaCl2. — In  section  442  it 
was  pointed  out  that  H2C03 
does  not  precipitate  CaC03,  as 
we  might  expect  according  to 
the  following  hypothetical 


•NNaCI 


CaCO, 


CaCO, 


FIG.  72 


equation: 


CaCl2+H2C03  ±*  CaC03|+  2HC1. 

The  reason  is  as  follows:  carbonic  acid  H2C03  is  a  very  weak 
acid  and  in  consequence  yields  but  very  few  C03~~  ions;  and 
although  CaCl2  gives  an  abundance  of  Ca++  ions,  the  concen- 
tration of  C03~"  ions  is  so  small  that  the  concentration  of 
CaC03  molecules  formed  is  less  than  the  M.S.  of  this  substance. 
Therefore  no  precipitation  of  CaC03  takes  place.  The  differ- 
ence in  behavior  of  H2CO3  and  Na2C03  toward  a  solution  of 
CaCl2  is  wholly  due  to  the  difference  in  their  tendencies  to  ionize, 
in  consequence  of  which  a  solution  of  H2CO3  contains  exceedingly 
few  C03  ions  as  compared  with  a  solution  of  Na2C03. 

The  behavior  of  H2C03  is  typical  of  that  of  all  weak  (little 
ionized)  electrolytes.  In  the  precipitation  of  salts,  weak  acids 
and  bases  are,  in  general,  less  efficient  precipitants  than  their 
salts,  since  the  latter  are  highly  ionized. 

450.  The  Precipitation  of  Magnesium  Hydroxide,  Mg(OH)2. 
— We  shall  next  discuss  in  detail  the  precipitation  of  Mg(OH)2, 


Applications  of  the  Ionic  Hypothesis  281 

not  so  much  because  of  the  chemical  importance  of  this  sub- 
stance, but  because  the  reactions  illustrate  in  a  striking  way  some 
of  the  most  important  principles  of  ionic  equilibrium. 

If  we  add  NaOH  to  a  solution  of  magnesium  chloride, 
MgCl2,  we  obtain  an  abundant  white  precipitate  of  Mg(OH)2, 
formed  as  follows: 

MgCl3+2NaOH^Mg(OH)a|+2NaCl. 

If  we  use  NH4OH  instead  of  NaOH  the  reaction  is  similar  but 
reaches  a  state  of  equilibrium  when  only  a  part  of  the  magnesium 
is  precipitated  : 


MgCl2+  2NH4OH^Mg(OH)2^+  2NH4C1. 

If  we  add  to  a  MgCl2  solution  a  solution  of  NH4OH  mixed  with 
sufficient  NH4C1,  no  precipitation  occurs.  We  shall  now  explain 
these  facts.  In  the  first  place  we  may  say  that  the  action  of 
NaOH  on  MgCl2  is  analogous  to  the  action  of  Na2CO3  on  CaCl2, 

the  two  equations  being 

* 

MgCl2+2NaOH^Mg(OH)2^-f-2NaCl, 
CaCl2+  NaCO3  ±*  CaC03^+  2NaCl. 


Both  MgCl2  and  NaOH,  like  CaCl2  and  Na2C03,  are  easily  solu- 
ble and  highly  ionized;  sodium  chloride,  one  of  the  products  in 
both  reactions,  is  also  highly  ionized.  The  other  product  in  the 
first  reaction,  Mg(OH)2,  is  but  slightly  soluble,  and  like  CaC03 
it  is  therefore  precipitated  almost  completely. 

If,  however,  we  use  NH4OH  instead  of  NaOH,  the  reaction  is 
far  from  complete.  The  reason  can  best  be  seen  by  the  aid  of 
Figs.  73  and  74.  Figure  73  shows  the  condition  of  the  solutions 
before  they  are  mixed;  Fig.  74  represents  the  condition  of  the 
mixture. 

It  will  be  recalled,  as  shown  in  Fig.  73,  that  NH4OH  is  but 
little  ionized.  Still  its  solution  yields  sufficient  OH~  ions  to 
produce  in  reaction  with  magnesium  chloride  solution  more 
Mg(OH)2  than  the  M.S.  of  the  latter  difficultly  soluble  substance. 
The  excess  of  Mg(OH)2  above  its  M.S.  precipitates,  Fig.  74.  As 
these  changes  go  on,  molecules  of  NH4OH  continue  to  ionize,  thus 


282 


Introduction  to  General  Chemistry 


bringing  into  the  solution  far  more  NH4+  ions  than  were 
originally  present  (cf.  Figs.  73  and  74).  The  presence  of  the 
large  excess  of  NH4+  ions  restricts  the  number  of  OH"  ions  to 
such  an  extent  that  a  state  of  equilibrium  is  reached  in  reaction, 
Mg(OH)2(dissolved)±5Mg+++2OH-, 

while  there  is  still  a  considerable  amount  of  magnesium  in  the 
form  of  Mg++  ions  and  MgCl2  molecules  left  in  the  solution. 
After  this  condition  is  reached 
no  more  Mg(OH)2  precipi- 
tates. We  therefore  conclude 
that  NH4OH  precipitates 
Mg(OH)2  only  partially,  (i) 
because  NH4OH  is  a  weak  or 
little  ionized  base,  and  (2) 


MgCI* 


NH,OH 


OH' 


U 


t — 


Mg(OH), 
I 


'  1  Mg<OH), 

I 
I 


FIG.  73 


FIG.  74 


because  the  accumulation  of  NH4+  ions  (Fig.  74)  finally  restricts 
the  OH~  concentration  to  so  small  a  value  that  the  Mg++ 
and  OH~  ions  are  just  in  equilibrium  with  the  amount  of 
Mg(OH)2  corresponding  to  its  M.S.  It  will  now  be  easy  to 
understand  why  no  Mg(OH)2  is  precipitated  when  NH4OH, 
mixed  with  considerable  NH4C1,  is  added  to  a  MgCl2  solution. 

The  NH4C1  furnishes  at  once  such  an  excess  of  NH4+  ions 
that  the  OH~  concentration  is  decreased  to  so  small  a  value  that 
less  Mg(OH)2  is  formed  in  the  reaction 

Mg+++2OH-±*Mg(OH)3 

than  corresponds  to  its  M.S.    Therefore  no  precipitation  occurs. 
451.  The  Precipitation  of  Ferric  Hydroxide,  Fe(OH)3. — The 

reaction 

FeCl3+3NH4OH±>Fe(OH)3NH-3NH4Cl 


Applications  of  the  Ionic  Hypothesis  283 

takes  place  with  the  practically  complete  precipitation  of  brown 
Fe(OH)3,  which  is  almost  insoluble  in  water.  The  completeness 
of  precipitation  is  not  noticeably  affected  by  the  addition  of  much 
NH4C1.  There  are  two  reasons  for  the  difference  in  behavior  of 
Fe(OH)3  and  Mg(OH)2:  (i)  the  former  is  much  more  insoluble 
in  water  than  the  latter,  so  that  the  M.S.  of  the  latter  (although 
small)  is  perhaps  1,000  times  as  large  as  that  of  the  former; 
(2)  Mg(OH)2  is  a  rather  strong  (highly  ionized)  base,  while 
Fe(OH)3  is  a  very  weak  base.  Therefore  in  the  reactions 

Mg+++2OH-^Mg(OH)2  (dissolved), 
Fe++++3OH-^Fe(OH)3  (dissolved), 

for  equal  concentrations  of  Mg++,  Fe+++,  and  OH~  ions  far 
less  Mg(OH)2  is  formed  than  Fe(OH)3.  The  presence  of  NH4C1 
decreases  the  OH~  concentration  of  NH4OH,  but  not  sufficiently 
to  prevent  the  practically  complete  precipitation  of  Fe(OH)3 
because  of  the  weakness  of  the  latter  and  its  exceedingly  small 
M.S.  The  weaker  a  base  and  the  smaller  its  M.S.,  the  more  com- 
pletely is  it  precipitated  by  NHAOH,  and  the  less  its  precipitation 
is  hindered  by  the  presence  of  ammonium  salts. 

452.  Classification  of  Precipitations.  —  The  various  examples 
of  precipitation  just  studied  cover  the  important  principles  con- 
cerned We  may  now  cite  a  few  additional  examples  of  each  of 
these  classes  of  precipitation.  In  the  reaction 

KBr+NaClO3^KClO3sH-NaBr,  (447) 

all  four  substances  are  highly  ionized,  and  all  but  KC1O3  are  very 
soluble.  The  latter  is  partially  precipitated  because  it  is  formed 
in  excess  of  its  not  very  large  M.S.  The  following  reactions  are 
of  this  class  : 


CaCl2+  2NaClO3^  2NaCl|+  Ca(ClO3)2, 
Pb(N03)a+2NaCl±?PbCU+2NaNOj,  (167) 

CaCl2+Na2S04^CaSO4NH-2NaCl.  (153) 

In  the  first  reaction  saturated  solutions  are  required  to  give  a 
precipitate  of  NaCl. 


284  Introduction  to  General  CJiemistry 

In  the  second  example  studied, 

CaCl2+Na2CO3^CaCO3^+2NaCl,  (448) 

the  precipitate  CaCO3  has  an  extremely  small  M.S.  Its  pre- 
cipitation by  Na2CO3  is  almost  complete.  Other  reactions  of 
this  class  are: 

AgNO3+NaCl=AgCH+NaNO3,  (382) 

AgN03+KBr=AgBr|+KN03,  (257) 

Pb(N03)2+CuS04=PbS04|+Cu(N03)2,  (167) 

BaCl2+Na2SO4  =  BaSO4|+  2NaCl,  (164) 

MgCla+2NaOH=Mg(OH)4+2NaCl.  '    (450) 

In  the  third  example  (449)  it  was  shown  that  H2CO3  did  not 
give  with  CaCl2  a  precipitate  of  CaC03,  because  the  former  is 
very  little  ionized  and  therefore  yields  very  few  C03  ions.  The 
following  pairs  of  substances  also  fail  to  give  precipitates,  because 
in  each  case  of  the  weakness  of  the  acid  coupled  with  the  moderate 
solubility  of  the  salt,  that  might  be  precipitated  : 

CaCl2  and  H3PO4, 
FeCl2  and  H2S, 
AgNO3  and  HC2H3O2. 

On  the  other  hand  the  reactions 

(158) 


FeCl2+  (NH4)2S  =  FeS^+  2NH4C1, 
AgN03+NaC2H3O2  =  AgC2H3O2^+NaNO3, 

give  abundant  precipitates,  because,  instead  of  the  weak  acids, 
we  use  their  salts,  which  are  highly  ionized. 

The  fourth  example,  which  belongs  to  Class  II,  was  taken  up 
in  contrast  to  the  fifth  example,  which  is  typical  of  a  fourth  class 
of  precipitation  reactions.  The  fourth  and  fifth  examples  were  : 

MgCla+2NaOH±>Mg(OH)a>H-2NaCl,  (450) 

MgCl2+  2NH4OH^Mg(OH)2|+  2NH4C1.  (450) 

The  precipitation  of  Mg(OH)2  is  nearly  complete  in  the  first 
reaction  but  only  partial  in  the  second,  owing  to  the  moderate 
solubility  of  Mg(OH)2  and  the  little  ionization  of  NH4OH, 
especially  in  the  presence  of  its  salts. 


Applications  of  the  Ionic  Hypothesis  285 

The  following  reaction  of  a  weak  electrolyte  (H2S)  also  results 
in  partial  precipitation: 


In  this  case  the  precipitation  is  prevented  by  the  presence  of 
an  excess  of  HC1  or  other  strong  acid,  because  of  the  suppression 
of  the  ionization  of  the  H2S  by  the  H+  ions  of  the  strong  acid. 
The  sixth  example  dealt  with  the  precipitation  of  Fe(OH)3: 

FeCl3+3NH4OH^Fe(OH)3|+3NH4Cl.  (451) 

In  this  case  the  precipitate  is  so  insoluble  (M.S.  so  small)  and 
so  weak  (little  ionized)  that  it  is  practically  completely  pre- 
cipitated by  NH4OH  even  in  the  presence  of  NH4C1.  Although 
we  call  NH4OH  a  weak  base,  it  is  enormously  stronger  than 
Fe(OH)3,  even  when  mixed  with  much  NH4C1.  Other  reactions 
which  fall  into  this  class  are: 

AlCl3+3NH4OH±5Al(OH)3sH-3NH4Cl,  (174) 

CrCl3+3NH4OH^Cr(OHU+3NH4Cl,  (344) 


Excess  of  NH4C1  in  the  first  two  cases,  and  of  HC1  or  H2S04  in 
the  last  two  cases,  fails  to  prevent  practically  complete  precipita- 
tion. 

453.  Precipitation  by  Adding  a  Substance  Having  a  Common 
Ion.  —  We  have  learned  in  the  foregoing  chapter  (432)  that  if  we 
add  to  the  solution  of  an  electrolyte,  AB,  enough  of  another 
highly  ionized  electrolyte,  AC,  having  a  common  ion,  A,  to 
increase  the  concentration  of  the  common  ion,  the  degree  of 
ionization  of  the  first  substance  will  be  suppressed.  If  now  the 
substance  A  B  is  not  very  soluble,  the  suppression  of  its  ioniza- 
tion caused  by  adding  AC  may  increase  the  concentration  of 
the  A  B  molecules  to  such  an  extent  as  to  exceed  the  M.S.  of  AB. 
In  consequence  part  of  AB  will  separate  out  as  a  precipitate. 
For  example,  if  a  few  bubbles  of  HC1  gas  are  passed  into  a 
saturated  solution  of  NaCl,  a  precipitate  of  NaCl  is  formed.  A 
similar  result  is  also  produced  by  adding  a  little  concentrated 
HC1  to  a  saturated  salt  solution.  In  each  case  the  ionization 


286  Introduction  to  General  Chemistry 

of  the  salt  is  suppressed  by  reason  of  the  increase  in  concentra- 
tion of  the  Cl~  ions,  and  the  concentration  of  the  molecular  NaCl 
increased  beyond  the  M.S.  of  this  substance.  Salt  precipi- 
tates until  the  concentration  of  molecular  NaCl  falls  to  the 
value  corresponding  to  its  M.S.  Another  example  illustrating 
the  same  principle  is  found  in  the  precipitation  of  KC103  from 
its  saturated  solution  by  the  addition  of  a  saturated  solution 
of  either  KBr  or  NaClO3.  We  may  say  that  as  a  general  rule  the 
total  solubility  of  a  salt  (molecular  and  ionic)  is  diminished  by  the 
presence  in  the  solution  of  another  electrolyte  having  a  common  ion. 
454.  Conditions  Favoring  Precipitation. — In  the  reaction 

AB+CD=AD+CB 

precipitation  will  occur  if  one  of  the  products,  say  AD,  is  formed 
as  molecules  in  greater  concentration  than  its  molecular  solu- 
bility. In  brief,  if  the  M.S.  is  exceeded,  precipitation  will  occur. 
Now  the  smaller  the  M.S.  of  AD,  the  more  probably  will  it  be 
exceeded. 

On  the  other  hand  the  M.S.  is  the  more  likely  to  be  exceeded 
the  greater  the  concentration  of  AD  which  tends  to  be  produced 
in  the  reaction.  The  various  factors  which  determine  the 
amount  of  AD  produced  (when  AD  is  soluble)  have  been  dis- 
cussed at  length  in  chapter  xviii  (Summary,  ,441). 

These  applications  of  the  ionic  hypothesis  have  the  following 
bearing  on  the  practice  of  precipitation.  In  the  first  place,  if 
we  are  to  precipitate  from  solution  an  " insoluble"  salt  of  a  weak 
acid  we  use  as  the  precipitant  a  soluble  salt  of  the  weak  acid 
instead  of  the  acid  itself,  since  the  former  is  highly  ionized, 
while  the  latter  is  not.  (The  term  precipitant  means  the  reagent 
added  to  cause  precipitation.)  If,  however,  we  are  to  precipitate 
an  insoluble  chloride  we  may  use  either  a  soluble  chloride  or 
hydrochloric  acid,  since  this  strong  acid  is  as  highly  ionized  as 
its  salt.  When  the  precipitant  is  added  to  a  given  solution,  a 
precipitate  may  not  appear  until  considerable  reagent  has  been 
added.  When  it  is  no  longer  possible  to  see  if  more  precipitate 
is  forming  with  further  additions  of  the  reagent,  a  small  portion 
of  the  mixture  is  filtered,  or,  better,  the  precipitate  is  allowed 


Applications  of  the  Ionic  Hypothesis  287 

to  settle,  and  the  clear  solution  is  tested  with  more  of  the  pre- 
cipitant. Only  moderate  excesses  of  the  precipitant  are  used 
as  a  rule,  since  in  many  cases  the  precipitant  reacts  farther  with 
the  precipitate  to  form  new  and  soluble  compounds,  with  the 
result  that  the  precipitate  dissolves  in  an  excess  of  the  reagent 
added. 

455.  Dissolving  Solid  Substances.  —  Substances  which  are 
not  readily  soluble  in  water  often  dissolve  easily  in  solutions  of 
other  electrolytes.  In  such  cases  we  may  imagine  that  chemical 
reaction  gives  rise  to  new  products  which  are  soluble  in  water. 
Here  is  a  case  in  point:  Calcium  hydroxide,  Ca(OH)2,  is  but 
slightly  soluble  in  water  (o.i2g.  in  looc.c.),  giving  a  very 
dilute  solution  known  as  limewater.  If  we  mix  a  few  grams  of 
Ca(OH)2  with  100  c.c.  of  water,  most  of  the  solid  remains  undis- 
solved.  If  now  we  add  dilute  HC1  to  the  mixture,  the  solid 
finally  passes  completely  into  solution.  The  explanation  is  as 
follows:  The  small  amount  of  dissolved  Ca(OH)2  (which  is  a 
strong  base)  is  neutralized  by  the  added  HC1  to  form  very  soluble 
CaCl2, 

Ca(OH)3+  2HC1^  CaCl2 


More  Ca(OH)2  then  dissolves  in  the  water  in  the  tendency  to 
keep  the  concentration  of  the  dissolved  Ca(OH)2  up  to  its  M.S.: 

Ca(OH)2±>Ca(OH)2^Ca+++20H- 
Solid        Dissolved 

As  fast  as  Ca(OH)2  passes  into  solution  it  reacts  with  the  HC1 
present.  If  the  chemically  equivalent  amount  of  HC1  is  added, 
all  Ca(OH)2  will  finally  dissolve,  and  the  solution  will  consist 
simply  of  CaCl2  dissolved  in  water. 

A  perfectly  analogous  reaction  is  found  in  the  dissolving  of 
the  difficultly  soluble,  strong  base  Mg(OH)2  in  dilute  HC1: 

Mg(OH)a+  2HC1  ±5  MgCla+  2H3O. 

Even  if  the  base  is  weak  -and  much  less  soluble  than  either 
Ca(OH)2  or  Mg(OH)2,  it  will  usually  dissolve  in  water  upon  the 
addition  of  a  strong  acid.  For  example,  Fe(OH)3  is  a  very  weak 


288  Introduction  to  General  Chemistry 

base  almost  insoluble  in  water;  it  dissolves  readily  in  dilute 
HC1,  forming  a  solution  of  ferric  chloride, 

Fe(OH)3+3HCl^FeCl3-f-3H20. 

The  stages  in  the  process  of  dissolving  may  be  considered  to  be 
comparable  to  those  in  the  case  of  the  dissolving  of  Ca(OH)2  in 
dilute  HC1. 

Most  bases,  with  the  exception  of  the  hydroxides  of  sodium, 
potassium,  ammonium,  and  barium,  are  very  little  soluble 
in  water.  All  such  so-called  insoluble  bases  dissolve  in  dilute 
HC1,  HNO3,  and  H2S04  to  form  clear  solutions,  if  their  cor- 
responding salts  with  these  acids  are  soluble  in  water. 

456.  Dissolving  Little  Soluble  Salts  of  Weak  Acids  by  Solu- 
tions of  Strong  Acids. — Silver  acetate  is  a  rather  difficultly  solu- 
ble salt  (i.og.  dissolves  in  100  c.c.  H20  at  18°)  which  is  easily 
made  by  precipitating  AgNO3  with  NaC2H3O2, 

AgN03+NaC2H302^AgC2H302+NaN03.  (452) 

If  we  mix  3  or  4  g.  of  AgC2H3O2  with  100  c.c.  of  H2O,  only  a  small 
portion  dissolves;  but  upon  addition  of  dilute  HNO3  the  whole 
of  the  solid  passes  into  solution.  Silver  acetate  is  the  salt  of 
the  weak  acid  HC2H3O2,  and,  as  we  have  already  learned  (428), 
a  strong  acid  reacts  more  or  less  completely  with  the  (soluble) 
salt  of  a  weak  acid  to  form  the  weak  acid  and  the  salt  of  the  strong 
acid.  This  was  shown  earlier  in  the  case  of  the  reaction 

HCl+NaC2H3O2±*HC2H302+NaCl.  (424) 

Nitric  acid  reacts  similarly  with  the  dissolved  portion  of  the 
AgC2H302, 

HNO3+AgC2H3O2^HC2H3O2+AgNO3. 

The  reaction  is  nearly  complete,  and  both  products  are  easily 
soluble.  The  dissolved  molecular  AgC2H3O2  being  thus  removed 
from  the  solution,  more  of  the  solid  passes  into  solution  in  the 
tendency  to  keep  the  concentration  of  molecular  AgC2H3O2  up 
to  its  M.S.  But  as  this  salt  reacts  with  the  HN03  present  as 


Applications  of  the  Ionic  Hypothesis  289 

soon  as  it  comes  into  solution,  its  M.S.  is  never  reached,  so  that 
finally  all  of  the  solid  passes  into  solution.  The  solution  con- 
sists largely  of  AgNO3  and  its  ions,  together  with  molecular 
acetic  acid. 

In  many  other  cases  so-called  ' 'insoluble"  salts  of  weak  acids 
dissolve  in  solutions  of  strong  acids  like  HC1,  HN03,  and  H2SO4. 
The  following  reactions  are  of  this  type: 

FeS+2HCl=H2S+FeCl2, 

CaC03+  2HC1 = H2CO3+  CaCl2, 

Ca3(P04)2+6HCl= 2H3P04+3CaCl2. 

However,  not  all  " insoluble"  salts  of  weak  acids  dissolve  in 
strong  acids.  For  example,  CuS,  which  comes  down  as  a  black 
precipitate  when  H2S  is  passed  into  a  solution  of  a  copper  salt, 
and  is  therefore  a  salt  of  the  very  weak  acid  H2S,  does  not  dissolve 
appreciably  in  cold  HC1.  The  reason  for  this  is  directly  trace- 
able to  the  extremely  small  M.S.  of  CuS.  In  general  the  smaller 
the  M.S.  of  a  salt  of  a  weak  acid  the  less  soluble  it  is  in  a  strong 
acid.  Other  examples  of  this  sort  are  found  in  Ag2S  and  HgS, 
neither  of  which  is  dissolved  appreciably  by  dilute  HC1  or 
H2S04. 

457.  Weak  Acids  and  Salts  of  Strong  Acids.— We  have 
already  learned  (282)  that  the  equilibrium  mixture  has  the  same 
composition  whether  we  start  with  one  pair  of  substances  of  a 
reaction  or  the  equivalent  amounts  of  the  other  pair.  In  accord 
with  this  principle  we  always  find  that  if  a  reaction  takes  place 
practically  completely  in  one  direction,  the  reverse  of  the  reaction 
does  not  succeed  under  the  same  conditions  of  temperature  and 
concentration.  In  sections  449  and  452  it  was  stated  that 
mixtures  of  the  following  pairs  of  substances  fail  to  give  pre- 
cipitates, although  little  soluble  salts  would  be  formed  by 
double  decomposition: 

H2CO3  and  CaCl2, 
H3PO4  and  CaCl2, 
H2S  and  FeCl2, 
HC2H3O2  and  AgNO3. 


290  Introduction  to  General  Chemistry 

Therefore  we  may  be  certain  that  calcium  carbonate,  calcium 
phosphate,  and  ferrous  sulphide  are  soluble  in  hydrochloric  acid, 
and  that  silver  acetate  is  soluble  in  nitric  acid.  Since  the  deter- 
mining factor  in  dissolving  each  of  these  salts  is  the  formation 
of  the  weak  acid,  we  may  go  farther  and  predict  that  any  strong 
acid  will  dissolve  these  salts.  Sometimes  a  new  insoluble  salt 
is  formed  by  the  strong  acid,  as  when  hydrochloric  acid  acts  on 
silver  acetate;  but  such  reactions  are  secondary  to  the  solution 
of  the  original  salts. 

458.  "Insoluble"  Salts  of  Strong  Acids.— The  "insoluble" 
salts  of  strong  acids  are  not  as  a  rule  dissolved  to  an  appre- 
ciable extent  by  solutions  of  other  strong  acids.    For  example, 
AgCl  is  not  appreciably  dissolved  by  HNO3,   although  the 
products  HC1  and  AgNO3  of  the  hypothetical  reaction 

AgCl+HNO3^HCl+AgNO3 

are  both  easily  soluble  substances.  A  reaction  in  which  both 
of  the  products  are  highly  ionized,  as  in  this  case,  falls  in  Class  I 
(414).  In  all  such  reactions  very  little  chemical  change  occurs, 
and  this  is  more  strikingly  true  the  more  dilute  the  solution.  As 
we  are  now  considering  the  case  where  one  of  the  substances 
taken  is  nearly  insoluble  in  water,  the  solution  of  this  substance 
must  be  exceedingly  dilute.  Comparing  the  action  of  HN03 
on  AgC2H3O2  and  AgCl,  we  may  say  that  the  first  reaction  takes 
place  readily  because  of  the  tendency  of  H+  and  C2H3O2~  to 
unite  nearly  completely  to  form  little  ionized  HC2H302;  and 
that  the  second  reaction  does  not  progress  far  because  of  the 
slight  tendency  for  H+  and  Cl"  ions  to  unite,  since  HC1  is 
nearly  completely  ionized  in  very  dilute  solution. 

459.  Evolution  of  a  Gas. — Substances  may  separate  from 
solutions  in  two  ways:  (i)  as  solid  precipitates  and  (2)  as  gases. 
We  have  considered  the  first  case  and  shall  now  take  up  the 
second,  and  we  shall  see  that  if  a  substance  separates  from  a 
solution  as  a  gas  the  effect  on  the  ionic  equilibrium  is  the  same 
as  if  the  substance  separated  as  a  solid.    The  principles  that 
apply  to  precipitation  apply  also,  with  slight  obvious  modifica- 
tions, to  gas  evolution.     Gases  have  limited  solubilities,  and 


Applications  of  the  Ionic  Hypothesis  291 

instead  of  the  M.S.  of  the  precipitate  we  have  the  molecular 
solubility  (M.S.)  of  the  gas.  Let  us  now  consider  a  few  well- 
known  reactions  as  illustrations. 

460.  The  Action  of  H2SO4  on  NaCl. — If  we  mix  dilute  solu- 
tions of  H2S04  and  NaCl  no  visible  effect  is  produced,  although 
in  the  solution  the  reaction 

H,S04+NaCl^HCl+NaHSO4 

takes  place  partially.  This  is  a  Class  I  reaction  (414)  since  all 
four  substances  are  easily  soluble  and  highly  ionized.  Therefore 
the  dilute  solution  contains  chiefly  the  ions  and  relatively  few 
molecules.  Nevertheless  some  HC1  molecules  are  formed  even 
in  dilute  solution,  but,  as  HC1  is  a  very  soluble  gas,  little  of  it 
escapes  from  the  solution  (251). 

On  the  other  hand  the  results  are  quite  different  if  but  little 
water  is  present.  In  making  hydrochloric  acid  (103)  58  g.  of 
NaCl,  100  g.  of  concentrated  H2S04,  and  30  g.  of  water  were 
mixed  in  a  flask  and  gently  heated. 

The  proportions  of  NaCl  and  H2S04  taken  were  those  indi- 
cated by  the  foregoing  equation,  since  the  100  g.  of  concentrated 
acid  taken  consists  of  98  g.  of  actual  H2SO4  and  2  g.  of  H2O. 
If  the  reaction  should  take  place  completely,  36. 5  g.  of  HC1 
would  be  formed.  This  is  far  more  HC1  than  the  32  g.  of  water 
present  can  hold  in  solution,  especially  when  the  mixture  is 
heated.  Therefore  HC1  gas  escapes  from  the  solution.  The  loss 
of  HC1  by  the  solution  impedes  the  reverse  action  on  the 
NaHSO4  present  and  so  causes  a  great  shift  to  the  right  of  the 
equilibrium  that  would  otherwise  be  reached  in  the  reaction 

H2SO4+NaCl±;HClt+NaHS04. 

As  a  consequence  this  reaction  goes  nearly  completely  from  left 
to  right  under  the  conditions  described  (103),  the  HC1  being 
given  off  as  gas.  In  equations  for  reactions  involving  gas 
evolution  the  gas  will  be  indicated  by  an  upward-pointing  arrow. 

461.  The  Action  of  HC1  on  CaCO3.— We  have  already  seen 
that  carbonic  acid,  H2C03,  does  not  precipitate  CaCO3  from  a 
CaCla  solution,  and  have  learned  that  this  is  because  H2CO3  is 


292  Introduction  to  General  Chemistry 

so  little  ionized  that  insufficient  molecular  CaCO3  is  formed  to 
exceed  its  M.S.     This  fact  indicates  that  the  reactions 

2HC1+  CaCO3  *=>  CaCl2+H2CO3, 


will  take  place  practically  completely,  since  in  all  reactions 
between  electrolytes  exactly  the  same  proportions  of  the  same 
products  result,  whether  we  start  with  one  pair  of  substances 
or  the  chemically  equivalent  amounts  of  the  other  pair  (282). 
The  dissolving  of  CaC03  in  dilute  HC1  takes  place  as  follows: 
CaCO3  first  dissolves  to  the  limit  of  its  (very  small)  M.S.  in  the 
water  present;  the  dissolved  molecules  then  ionize  rather  highly: 

CaCO3±?Ca+++C03~. 

The  CO3~"  ions  unite  nearly  completely  with  the  H+  ions  of 
the  highly  ionized  HC1  present, 


and  to  a  small  extent  Ca++  and  Cl~  ions  unite  to  form  (easily 
ionized)  CaCl2.  The  nearly  complete  removal  of  CO3"  ions 
allows  the  further  ionization  of  CaCO3,  and  this  change  permits 
the  passage  of  more  solid  CaC03  into  solution.  The  quantitative 
relations  are  such  that  these  changes  continue  until  all  CaC03 
has  dissolved.  Incidentally  the  H2C03,  which  is  unstable,  dis- 
sociates, to  a  large  extent,  into  water  and  CO2, 


and,  as  the  latter  is  not  very  soluble,  much  of  it  passes  off  as 
a  gas. 

The  several  reactions  are  shown  in  the  following  diagram  : 

CO3  (gas) 

it   ' 
H2CO3^H2O+CO2 

It 


CaCO3  ±5  CaCO3  ±*  Ca+++  CO3-  - 
(Solid)  It 

CaCl2 


Applications  of  the  Ionic  Hypothesis  293 

462.  The  Action  of  NaOH  on  NH4C1.  —  Another  example  of 
gas  evolution,  which,  however,  does  not  introduce  any  new 
principle,  is  found  in  the  reaction 

NaOH+NH4Cl  =  NaCl+NH4OH, 

which  takes  place  more  or  less  completely  when  solutions  of  the 
two  initial  substances  are  mixed.  This  reaction  was  studied 
under  Class  II  (426)  ,  where  it  was  pointed  out  that  it  takes  place 
nearly  completely  because  NH4OH  is  a  weak  base.  This  base 
is  also  unstable,  readily  dissociating,  thus, 


and,  since  NH3  is  a  gas,  it  will  in  part  escape  from  the  solution. 
The  more  concentrated  the  solution  and  the  higher  the  tempera- 
ture the  more  completely  will  the  NH3  be  evolved  as  gas.  The 
loss  of  NH3  from  the  solution  promotes  the  dissociation  of 
NH4OH,  and  this  in  turn  favors  a  further  shift  in  equilibrium 
from  left  to  right  in  the  main  reaction.  The  various  reactions 
and  their  relations  are  fully  shown  in  the  following  diagram: 

NH3  (gas) 

It 
NH4OH±5H20+NH3 

It 
NH4C1^C1-+NH4+ 

NaOH^Na++OH- 

It 
NaCl 

463.  The  Factors  Governing  Gas  Evolution.—  The  various 
factors  which  are  favorable  to  gas  evolution  are  very  similar 
to  those  which  were  found  to  favor  precipitation,  although  there 
are  some  differences  aside  from  the  fact  that  in  the  one  case  we 
are  dealing  with  a  gas  and  in  the  other  with  a  solid  product.  If 
one  of  the  products  of  a  reaction  is  gaseous  it  will  be  given  off 
from  the  solution,  the  more  completely,  the  larger  the  propor- 
tion of  it  formed  in  the  reaction  and  the  less  soluble  it  is.  In 
these  respects  gas  evolution  is  completely  analogous  to  precipita- 
tion. Since  all  gases  are  less  soluble  at  high  than  at  low  tempera- 
tures, gas  evolution  is  always  more  complete  the  higher  the 


2Q4  Introduction  to  General  Chemistry 

temperature.  Gas  evolution  and  precipitation  differ  in  one 
very  important  additional  respect:  at  a  given  temperature  the 
M.S.  of  a  precipitate  has  a  fixed  value,  while  that  of  a  gas  depends 
upon  the  pressure  of  the  gas  above  its  solution.  In  most  cases 
the  total  solubility  and  therefore  also  M.S.  of  a  gas  is  directly 
proportional  to  its  (partial)  pressure  (Henry's  Law,  126).  If 
during  gas  evolution  the  partial  pressure  of  the  gas  given  off 
is  kept  down  by  removing  the  gas  (as  by  blowing  it  away  with  a 
stream  of  air)  as  fast  as  it  is  liberated,  its  M.S.  will  be  reduced  to 
a  vanishingly  small  value.  In  consequence  the  dissolved  gas 
will  be  practically  or  even  completely  removed  from  the  solution. 
Thus,  in  the  reaction  between  NaOH  and  NH4C1,  if  a  current  of 
air  is  blown  through  the  solution  every  trace  of  NH3  will  finally 
be  removed,  so  that  the  reaction  will  be  absolutely  complete. 
The  remaining  solution  will  contain  only  common  salt.  The 
same  result  is  attained  if  the  solution  is  boiled,  in  which  case  the 
evolved  steam  takes  the  place  of  the  air  current.  The  high 
temperature  also  hastens  the  removal  of  the  NH3.  All  reactions 
giving  gases  which  follow  Henry's  Law  may  be  driven  to  com- 
pletion by  the  continuous  removal  of  the  gas  by  means  of  a 
current  of  an  inert  gas  or  by  steam  produced  when  the  solution 
is  boiled.  We  have  seen  (455,  461)  that  the  reason  why  a  little 
soluble  salt  dissolves  is  the  efficient  removal  from  solution  of  one 
or  both  of  its  ions  to  form  some  new  substance,  which  of  course 
must  be  soluble  or  volatile  if  the  resulting  mixture  is  to  be  a 
solution.  In  the  reactions  studied  so  far  the  removal  of  ions 
has  been  accomplished  by  the  formation  of  little  ionized  or 
little  soluble  substances.  There  are  other  ways  of  removing 
ions.  These  we  shall  take  up  later.  We  shall  find  that  the 
dissolving  of  little  soluble  substances  in  question  depend  upon 
the  same  fundamental  principle,  and  that  these  new  cases  differ 
from  those  studied  in  this  chapter  only  in  secondary  ways,  the 
means  by  which  the  ions  in  question  are  removed  from  the  solu- 
tion (532,  560,  626). 

464.  The  Value  of  the  Ionic  Hypothesis. — In  chapters  xviii 
and  xix,  we  have  applied  the  ionic  hypothesis  to  the  interpreta- 
tion of  reactions  between  acids,  bases,  and  salts  and  have  seen 


Applications  of  the  Ionic  Hypothesis  295 

that  this  hypothesis  leads  to  fairly  simple  explanations  of  a 
great  variety  of  facts.  Furthermore  we  have  seen  that  if  we 
know  the  degrees  of  ionization  and  the  solubilities  of  the  sub- 
stances concerned  in  any  reaction  we  are  able  to  predict  what 
the  result  of  the  reaction  will  be.  Herein  lies  the  enormous 
practical  value  of  the  ionic  hypothesis. 

In  chapter  xvii  (412)  we  called  attention  to  some  of  the 
glaring  inconsistencies  of  the  hypothesis;  but  we  have  also 
pointed  out  that  the  practical  value  of  any  hypothesis  is  not  its 
truth  but  its  usefulness.  Having  now,  we  hope,  shown  its  use- 
fulness, we  shall  in  later  chapters  consider  whether  it  is  true 
(chaps,  xx,  xxvii). 


CHAPTER  XX 
ELECTROCHEMISTRY 

465.  Introduction. — The  present  chapter  will  deal  first  with 
some  of  the  marvelous  developments  of  our  knowledge  of  elec- 
tricity and  matter  during  the  last  two  decades.    We  now  have 
good  reason  for  believing  that  electricity  like  matter  is  of  a 
granular  or  " atomic'1  nature.    The  grains  or  "atoms"  of  free 
electricity  are  all  exactly  alike  and  of  the  variety  known  as 
negative  electricity.    These  grains  are  called  electrons.    Posi- 
tive electricity  is  not  known  in  a  free  state;   that  is,  it  is  only 
known  as  a  positive  charge  on  ordinary  chemical  atoms  or 
larger  masses  of  matter. 

466.  The  Granular  Nature  of  Electricity;    Electrons. — In 
chapter  xvii  (403)  it  was  shown  that  Faraday's  Law  of  Electro- 
chemical Equivalents  leads  directly  to  the  conclusion  that  all 
univqlent  ions,  in  solution,  carry  equal  charges  of  electricity  (404) . 
The  charge  on  a  single  univalent  ion  may  be  called  a  unit  charge. 
Each  bivalent  ion  has  two  unit  charges,  each  trivalent  ion  three 
unit  charges,  etc.    As  early  as  1874  Stoney  pointed  out  that  these 
facts  indicate  that  electricity  is  granular  in  nature,  that  each 
univalent  atom  is  associated  with  one  such  granule  to  form  a 
univalent  ion,  that  a  bivalent  ion  is  an  atom  with  two  granules 
of  electricity,  etc.     Furthermore  a  little  later  he  proposed  to 
call  the  quantity  of  electricity  of  a  single  granule  an  electron. 
This  quantity  is  exceedingly  minute.    The  common  unit  of 
quantity,  one  coulomb,  is  equal  to  more  than  a  billion  billion 
electrons.    According  to  present-day  usage  the  term  electron 
means  a  single  electronegative  granule  of  electricity. 

467.  Proof  of  the  Existence  of  Electrons. — Although  evidence 
was  gradually  accumulating  during  the  last  quarter  of  the  nine- 
teenth century,  it  was  not  until  more  recently  that  positive 
evidence  was  obtained  that  electricity  is  granular  and  is  made 
up  of  electrons.    The  crowning  work  was  that  of  Professor 

296 


Electrochemistry 


297 


Robert  Millikan,  an  American  physicist  who  showed  that  when 
a  very  small  sphere  is  charged  with  more  and  more  electricity 
the  quantity  of  electricity  increases  by  small,  equal  additions,  and 
not  continuously.  This  is  exactly  what  we  should  expect  if  the 
charge  is  made  up  of  a  small  number  of  electrons. 

The  spheres  used  were  oil  drops  of  microscopic  size,  not 
much  larger,  in  fact,  than  the  particles  of  dust  that  can  be  seen 
floating  in  the  air  when  a  beam  of  sunlight  penetrates  a  dark- 
ened room.  A  drop  was  made  visible  by  a  beam  of  bright  light 
and  was  viewed  through  a  short-focus  telescope.  In  still,  dust- 
free  air  the  drop  fell,  under  the  action  of  gravity,  at  a  constant 
velocity  that  could  be  accurately  measured.  It  is  interesting 


FIG.  75 

to  note  that  although  every  body,  however  small,  will  fall  with 
steadily  increasing  velocity  in  a  vacuum,  a  very  small  body 
falls  with  constant  velocity  in  air,  owing  to  the  viscosity  of  the 
latter.  The  drops  used  fell  with  a  velocity  of  about  one  milli- 
meter per  second.  The  principal  parts  of  Millikan' s  apparatus 
are  shown  in  Fig.  75 :  M  and  N  are  parallel  metal  plates  insulated 
from  one  another  and  connected  through  a  switch  to  the  terminals 
of  a  high-potential  battery,  B.  The  upper  plate,  M,  can  be 
charged  positively  and  the  lower  one,  N,  negatively.  A  minute 
oil  drop,  D,  is  caused  to  fall  into  the  space  between  M  and  N 
through  a  pinhole  in  the  cjnter  of  M,  and  its  rate  of  fall  is 
measured  while  M  and  N  are  uncharged,  observations  being 
made  with  a  telescope,  T.  A  minute  negative  charge  is  then 
given  to  the  drop  (in  a  way  that  need  not  be  considered  at 
present),  and  the  plates  M  and  N  are  charged.  The  drop  is 


298  Introduction  to  General  Chemistry 

now  attracted  by  M  and  repelled  by  N,  so  that  it  moves  upward. 
When  it  is  close  to  M  the  electric  field  is  switched  off  (S,  switch; 
E,  earth),  and  the  drop  is  again  allowed  to  fall,  and  its  speed 
(time  of  fall)  is  again  measured.  With  uncharged  plates  (field 
off)  the  drop  falls  at  exactly  the  same  rate,  whether  it  is  charged 
or  uncharged.  Next  its  speed  upward  is  measured  with  the 
field  on.  This  speed  is  always  the  same  as  long  as  the  charges 
on  M  and  N  remain  constant  (constant  field),  and  the  charge 
on  the  drop  is  unchanged.  But  increase  of  negative  charge  on 
the  drop  increases  the  upward  speed,  and  decrease  of  negative 
charge  decreases  the  upward  speed,  the  field  remaining  constant. 
The  speed  upward  is  a  measure  of  the  force  of  electrical  attrac- 
tion by  M  and  repulsion  by  N  of  the  charged  drop  and  is  there- 
fore a  measure  of  the  charge  on  the  drop.  A  drop  could  be  made 
to  make  hundreds  of  trips  up  and  down.  The  downward 
velocity  (field  off)  was  always  the  same;  the  upward  velocity 
(field  on)  varied  with  the  charge.  The  charge  for  each  upward 
speed  was  found  by  a  simple  calculation.1  It  turned  out  that 
the  charge  on  the  drop  was  in  every  case  a  multiple  by  a  whole 
number  of  the  smallest  possible  charge  on  the  drop.  Thousands 
of  observations  were  made  in  these  experiments,  but  not  an 
exception  was  found  to  the  foregoing  statement.  This  proves 
conclusively  that  electricity  is  granular  in  nature.  It  has  been 
shown  in  other  ways  that  the  granules  of  electricity  composing 
the  charge  on  an  oil  drop  are  of  the  same  magnitude  as  the  unit 
charges  of  ions  of  electrolytes  in  solution.  We  may  therefore  call 
them  electrons  and  say  that  all  electric  charges  are  made  up  of 
one  or  more  electrons.  In  Millikan's  experiments  the  oil  drops 
used  were  observed  to  carry  all  possible  charges  from  a  single 
electron  to  over  a  hundred  electrons;  in  no  single  case  was  a 
fraction  of  an  electron  found.  The  electron  is  therefore  the 
smallest  indivisible  particle  of  electricity. 

468.  The  Nature  of  an  Electric  Current.— The  relation  of 
an  electric  charge  to  an  electric  current  was  first  clearly  estab- 
lished in  1876  by  the  American  physicist  Rowland,  who  showed 

XA  popular  account  of  Professor  Millikan's  work  is  given  in  his  book,  The 
Electron.  Chicago:  The  University  of  Chicago  Press,  1917. 


Electrochemistry  299 

that  when  an  electrically  charged  gilt  disk  was  very  rapidly 
rotated  it  produced  the  same  sort  of  deflection  on  a  magnetic 
needle  as  that  due  to  a  current  of  electricity  flowing  through  a 
wire  having  the  same  position  with  respect  to  the  needle  as 
that  occupied  by  the  disk,  Fig.  76.  This  experiment  proved 
that  a  current  of  electricity  is  nothing  but  an  electric  charge 
in  motion,  just  as  a  current  of  water  is  nothing  but  water  in 
motion. 

469.  The   Electron   Theory   of   Electric    Currents. — If   we 
accept  the  view  that  an  electric  current  is  an  electric  charge  in 
motion,  and  also  take  into  account  the  fact  that  an  electric 
charge  is  merely  an  assemblage  of  electrons,  we  are  at  once  led 
to  the  supposition  that  a  current  in  a  wire 

is  only  a  stream  of  electrons  passing  through 
the  wire. 

470.  The  Structure  of  an  Atom. — If 
we  think  of  a  metal  wire  as  made  up  of 
" solid,"   impenetrable   atoms,   it   is   not 
very    reasonable    to   imagine    that   par- 
ticles of  electricity  (electrons)  could  pass 

through     it.    However,    physicists    and  FlG    6 

chemists  have  in  recent  times  come  to 
the  conclusion  that  an  atom  is  by  no  means  a  homogeneous, 
solid  lump,  but  that  it  is  a  rather  complex  structure,  con- 
sisting largely  or  even  wholly  of  negative  electrons  rotating 
in  more  or  less  circular  orbits  about  a  positively  charged 
nucleus.  The  sum  of  the  negative  charges  of  the  electrons 
is  exactly  equal  to  the  positive  charge  of  the  nucleus,  so  that, 
as  a  whole,  an  atom  has  no  excess  of  either  kind  of  elec- 
tricity. The  structure  of  an  atom  may  be  likened  to  that  of 
the  solar  system,  in  which  the  sun  corresponds  to  the  nucleus 
and  the  planets  to  the  surrounding  electrons.  The  distances 
between  the  electrons  composing  an  atom  are  probably  large 
compared  with  the  size  of  an  electron,  so  that  a  stray  electron 
might  pass  through  an  atom  with  the  same  ease  that  a  comet 
passes  through  the  solar  system,  or  a  bullet  may  pass  through 
a  flock  of  birds  without  striking  any  one  of  them. 


300  Introduction  to  General  Chemistry 

471.  How  a  Wire  "Carries  a  Current." — If  we  think  of  a 
wire  as  made  up  of  atoms  of  the  sort  here  pictured,  it  is  easy  to 
see  how  a  stream  of  electrons  might  pass  through  it.     In  a  wire 
(not  connected  with  any  electrical  source)  some  electrons  are 
continuously  becoming  detached  from   their  original  atoms; 
these  probably  move  through  and  among  the  atoms,  occasionally 
replacing,  for  the  time  being,  those  that  have  been  lost  by  other 
atoms.    A  metal  always  contains  more  or  less  of  these  free, 
wandering  electrons,  as  well  as  an  equal  number  of  atoms  which 
are  deficient  in  electrons.    When  the  terminals  of  a  battery  are 
joined  by  a  wire,  the  positive  pole  of  the  battery  attracts  and 
the  negative  pole  repels  the  free  electrons  of  the  wire.    This 
causes  a  drift  of  electrons  along  the  wire,  and  this  drift  of  electrons 
constitutes  the  current  in  the  wire.    The  progress  of  electrons  in 
the  direction  of  the  drift  is  slow,  a  matter  of  a  few  centimeters 
per  minute.    The  well-known  fact  that  the  effect  of  closing  an 
electric  circuit  is  felt  almost  instantaneously  at  a  great  distance 
(as  illustrated,  for  example,  by  our  everyday  telephone  experi- 
ence) is  explained  by  the  assumption  that  all  the  mobile  electrons 
in  the  wires  of  the  circuit  move  forward  at  the  instant  the  circuit 
is  closed.    The  case  is  just  like  that  of  drawing  water  from  a 
supply  pipe;  water  flows  out  the  instant  the  faucet  is  opened, 
being  pushed  forward  along  the  whole  pipe  by  the  water  forced 
into  the  pipe  by  the  pump. 

472.  The    Direction   of   an   Electric    Current. — Before   the 
nature  of  an  electric  current  had  been  discovered  it  was  cus- 
tomary to  consider  that  the  current  in  the  wire  flowed  from  the 
positive  to  the  negative  pole.     Since  the  drift  or  flow  of  electrons 
is  in  the  opposite  direction,  there  is  danger  of  misunderstanding 
in  speaking  of  the  direction  of  the  current.     It  is  perhaps  best 
to  speak  of  the  direction  of  the  negative  current,  which  is  then 
the  direction  of  drift  of  the  electrons. 

473.  Nonconductors  of  Electricity. — All  metals  are  good  con- 
ductors, but  the  non-metals  are  practically  nonconductors  or  in- 
sulators.   To  account  for  this  difference  we  have  only  to  suppose 
that  a  non-metal,  like  sulfur,  contains  but  very  few  free  or  mobile 
electrons  and  therefore  has  very  little  ability  to  carry  a  current 


Electrochemistry  301 

474.  Production  of  Electric  Charges  by  Friction. — If  a  glass 
rod  is  rubbed  with  a  piece  of  silk,  the  former  takes  on  a  positive, 
the  latter  a  negative,  charge.     This  is  explained  by  assuming 
that  a  few  stray  electrons  of  the  glass  have  been  " wiped  off"  by 
the  silk.    The  rubbing  of  the  glass  by  the  silk  is  of  importance 
only  in  insuring  intimate  contact  between  the  two.    Another 
example  of  similar  nature  is  found  in  the  familiar  electrification 
of  the  hair  when  combed  with  a  hard-rubber  comb  in  dry  weather. 
Here  the  comb  acquires  a  negative  charge  and  the  hair  a  posi- 
tive one.    In  general,  when  any  two  different  substances  are  brought 
together  they  become  electrified  with  opposite  charges.    This  may  be 
taken  to  mean  that  electrons  accumulate  in  excess  more  easily 
on  some  kinds  of  matter  than  on  others. 

475.  Cathode    Rays.— 
When  a  high-voltage  electric 
current  is  passed  through  a 
Crookes  tube,  Fig.  77,  which 
is   an  evacuated  glass  bulb 
having    a    metallic  cathode, 
C,  and  an  anode,    A,  rays, 

known  as  cathode  rays,  are  given  off  by  the  cathode  and  cause 
a  yellowish-green  fluorescence  of  the  opposite  end  of  the  tube. 
These  rays  are  readily  stopped  by  a  sheet  of  metal,  as  shown  by 
the  fact  that  a  screen  (in  the  form  of  a  Maltese  cross)  casts  its 
shadow  on  the  glass.  Even  transparent  substances  like  glass 
do  not  transmit  the  cathode  rays  any  better  than  do  metal  sheets 
of  comparable  thickness.  Extremely  thin  sheets  of  material  like 
aluminum  or  gold  leaf  permit  partial  transmission  of  the  cathode 
rays. 

476.  X-Rays. — Cathode  rays  produce  X-rays,  also  known  as 
Roentgen  rays,  which  radiate  from  any  object  struck  by  the 
former.    A  modern  X-ray  tube  is  shown  in  Fig.  78.    This  is  a 
modified  Crookes  tube,  intended  for  the  use  of  large  currents 
and  the  production  of  powerful  X-rays.     The  cathode  rays  come 
from  the  specially  constructed  cathode,  C,  and  strike  a  target,  T, 
made  of  metallic  tungsten,  which  metal  is  chosen  because  of  its 
very  high  melting-point  (3000°).     When  the  cathode  rays  are 


302 


Introduction  to  General  Chemistry 


stopped  by  the  target,  part  of  their  energy  is  transformed  into 
X-rays,  and  the  balance  appears  as  heat,  so  that  the  target 
becomes  red,  or  even  white,  hot.  Recent  work  has  proved  that 
the  X-rays,  which  are  very  different  from  the  cathode  rays,  are, 
like  visible  light,  vibrations  of  the  so-called  luminous  ether  and 
differ  from  visible  light  in  having  wave-lengths  about  one- 
thousandth  as  great  as  the  latter. 

477.  The  Nature  of  Cathode  Rays. — The  extensive  investiga- 
tions of  Sir  William  Crookes  on  cathode  rays,  during  the  seventies 
of  the  last  century,  led  this  famous  English  physicist  to  con- 
clude that  these  rays  were  matter  in  a  highly  rarefied  or 


FIG.  78 

ultra-gaseous  state,  which  he  called  a  fourth  state  of  matter 
(the  other  three  states  being  solid,  liquid,  and  gaseous).  And 
in  the  light  of  our  present  knowledge  of  the  real  nature  of  these 
remarkable  rays  we  must  admit  that  Crookes's  conclusion  was 
substantially  correct,  although  it  was  by  no  means  the  last  word 
on  the  subject. 

It  has  long  been  known  that  cathode  rays  travel  in  a  straight 
line  in  a  vacuum,  but  that  they  may  be  deflected  in  an  arc  of  a 
circle  by  a  transverse  magnetic  field.  The  apparatus  shown  in 
Fig.  79  serves  for  lecture  demonstration  of  this  interesting 
phenomenon.  A  narrow  beam  of  rays  coming  from  the  cathode 
and  passing  through  a  slit  in  a  mica  plate  strikes  along  a  screen 
covered  with  a  specially  prepared  form  of  zinc  sulfide,  which 
becomes  luminous  in  the  line  where  it  is  struck  by  the  rays.  If 
now  a  horseshoe  magnet  is  presented  so  that  the  N  pole  is  above 


Electrochemistry  303 

the  plane  of  the  paper  and  the  S  pole  below  it,  the  beam  is 
deflected  to  the  position  of  the  curved  line. 

It  is  a  well-known  fact  that  a  wire,  free  to  move,  is  deflected 
by  a  magnetic  field  when  a  current  is  passed  through  it.  The 
direction  of  deflection  of  the  wire  is  determined  by  the  direction 
of  the  current  in  the  wire.  The  deflection  of  the  cathode  rays 
by  a  magnetic  field  indicates  that  the  rays  are  electricity  in 
motion,  the  direction  of  deflection  corresponding  to  that  of  a 
stream  of  negative  electricity  coming  from  the  cathode,  which  is, 
of  course,  the  negative  electrode.  If  we  grant  that  the  current 
in  the  wire  leading  to  the  cathode  is,  in  reality,  only  a  stream  of 
negative  electrons  in  the  wire,  we  have  only  to  suppose  that  these 


FIG.  79 

electrons  do  not  stop  on  reaching  the  cathode,  but  shoot  out  from  the 
surface  of  the  latter  and  thus  constitute  the  cathode  rays. 

478.  Proof  that  the  Cathode  Rays  Are  a  Stream  of  Electrons. 
— The  conclusive  proof  that  the  cathode  rays  are  a  stream  of 
negative  electricity  (presumably  electrons,  since  all  negative 
charges  consist  of  electrons)  was  furnished  by  the  work  of 
Perrin,  a  French  physicist.  Perrin's  apparatus  is  shown  in 
Fig.  80.  It  was  a  special  form  of  Crookes  tube  having  the 
cathode  at  C,  the  anode  at  A,  and  at  B  an  insulated  metal 
receiver,  into  which  the  cathode  rays  could  be  deflected  by  means 
of  a  magnet.  This  receiver  was  connected  by  a  wire  to  an  elec- 
troscope, capable  of  detecting  any  electric  charge  given  to  the 
box  and  determining  its  sign,  whether  positive  or  negative. 
When  the  cathode  rays  were  started  no  charge  passed  into  the 
receiver  until  the  rays  were  magnetically  deflected  so  as  to  fall 
into  the  receiver;  then  the  latter  acquired  a  large  negative 


304  Introduction  to  General  Chemistry 

charge.  To  guard  against  stray  electric  charges  the  receiver 
was  surrounded  by  a  metal  shield  connected  to  the  earth,  E. 
The  experiments  above  described,  together  with  many  other 
facts,  have  led  to  the  conclusion  that  the  cathode  rays  are  com- 
posed of  negative  electrons  shot  out  from  the  cathode  with  high 
velocity. 

479.  The  Mass  of  an  Electron. — An  electron  behaves  as 
thought  it  had  mass.  In  the  first  place  we  know  that  moving 
electrons  have  energy,  since  the  cathode  rays  can  produce  light, 
heat,  and  X-rays,  all  of  which  are  forms  of  energy.  Since  the 
kinetic  energy  of  a  moving  body  is  proportional  to  the  product 


FIG.  80 

of  its  mass  and  the  square  of  its  velocity,  we  can  account  for 
the  energy  of  the  cathode  rays  by  assuming  the  electrons  to 
have  mass.  Furthermore  the  fact  that  it  requires  an  appreciable 
magnetic  force  to  deflect  the  cathode  rays  and  that  the  extent 
of  the  deflection  (for  rays  of  a  given  velocity)  is  proportional 
to  the  strength  of  the  magnetic  force  is  also  evidence  that  elec- 
trons have  mass.  One  of  Newton's  laws  is  to  the  effect  that  a 
moving  mass  continues  in  a  straight  line  unless  acted  upon  by  a 
transverse  force.  Conversely,  if  a  force  is  required  to  deflect 
a  moving  electron,  we  are  warranted  in  assuming  that  the  latter 
has  mass.  '  By  methods  that  we  cannot  explain  here  it  has  been 
shown  that  the  mass  of  an  electron  is  about  one  eighteen-hundredth 
that  of  an  atom  of  hydrogen. 

480.  The  Beta  Rays  of  Radium. — The  spectacular  properties 
of  radium  have  been  brought  to  the  attention  of  nearly  everyone, 


Electrochemistry  305 

whether  he  is  a  student  of  chemistry  or  not.  Radium  gives  out 
three  kinds  of  rays,  the  alpha,  beta,  and  gamma  rays.  Of  these 
the  beta  rays  very  closely  resemble  the  cathode  rays.  Like 
cathode  rays  they  are  deflected  by  a  magnetic  field  in  a  direc- 
tion which  indicates  that  they  too  are  a  stream  of  electrons 
shot  out  with  high  velocity  from  the  radium.  Radium  is,  by 
all  ordinary  tests,  an  element.  It  resembles  barium  as  closely 
as  potassium  resembles  sodium.  Here  then  is  an  element  that 
spontaneously  gives  off  negative  electricity  in  the  form  of  elec- 
trons shot  out  with  great  velocity. 

The  alpha  rays  have  been  proved  to  be  atoms  of  the  element 
helium,  He  (atomic  weight  =  4),  each  of  which  carries  a  double 
positive  charge.  These  rays  are  also  shot  out  with  high  velocity. 
The  gamma  rays  are  identical  with  X-rays. 

481.  The    Disintegration    Hypothesis.— The    extraordinary 
behavior  of  radium  has  been  satisfactorily  explained  by  the 
disintegration  hypothesis   of   Rutherford   and   Soddy.    These 
scientists  assumed  that  a  radium  atom  is  not  a  homogeneous 
solid  particle  but  a  very  complex  structure  made  up  of  electrons 
revolving  rapidly  in  more  or  less  circular  orbits  about  a  nucleus 
of  positive  electricity  in  the  manner  already  described.     It  is 
further  assumed  that  an  atom  of  radium  may  become  unstable 
and  throw  of  a  single  electron  (beta  ray)  or  a  larger  mass  (an  atom 
of  helium ,  which  is  an  alpha  ray),  leaving  behind  an  atomic  residue 
of  smaller  mass  and  therefore  smaller  atomic  weight.    This  hy- 
pothesis is  in  complete  accord  with  all  known  facts  concerning 
radium  and  radioactive  phenomena. 

482.  The  Electrical  Nature  of  Matter. — The  study  of  radio- 
active substances,  of  which,  in  addition  to  radium,  about  thirty 
are  known,  has  led  to  the  conclusion  that  the  atoms  of  all  ele- 
ments, whether  radioactive  or  not,  are  constructed  on  about  the 
same  plan  as  that  of  radium.    According  to  this  hypothesis  the 
atom  of  one  element  differs  from  that  of  another  only  in  the  number 
and  arrangement  of  the  electrons  composing  it.    The  mass  of  an 
atom  is,  at  least  in  part,  accounted  for  by  the  mass  of  the  elec- 
trons composing  it.     All  matter  is  considered  to  be  of  electrical 
origin. 


306  Introduction  to  General  Chemistry 

483.  The  Nature  of  an  Ion. — A  single  sodium  ion  is  an  atom 
of  sodium  having  a  single  positive  charge  of  electricity  equal  in 
quantity  but  opposite  in  sign  to  that  of  an  electron.    The 
simplest  explanation  of  the  difference  between  an  ion  and  an 
atom  of  sodium  is  found  in  the  assumption  that  the  ion  is  an 
atom  which  has  lost  one  electron.*   The  atom  was  originally  elec- 
trically neutral,  because  the  positive  charge  of  its  nucleus  was 
just  equal  to  the  sum  of  the  negative  charges  of  its  surrounding 
electrons.    If  one  electron  is  lost,  the  atom  will  have  an  excess 
positive  charge  just  equal  in  magnitude  to  that  of  one  electron. 
Since  metallic  atoms  all  form  positive  ions  we  conclude  that  all 
such  atoms  are  able  to  lose  electrons.    Moreover,  an  atom  of 
a  univalent  metal  can  lose  but  one  electron  and  its  ion  will  have 
a  single  unit  charge,  thus, 

Na(atom)->Na++one  electron. 

A  bivalent  atom  can  lose  two  electrons, 

Ca(atom)->Ca+++two  electrons. 

A  trivalent  atom,  such  as  that  of  aluminum,  can  lose  three  elec- 
trons, etc. 

Later  work  has  shown  that  the  ions  are  undoubtedly  hydrated 
to  some  extent.    The  actual  formula  of  sodium  ion  might  be 

represented  thus: 

Na(H20);  +  . 

The  subscript  x  represents -a  small  integer,  probably  2  or  3.  In 
practice  we  do  not  include  the  water  in  the  formula,  since  in 
the  first  place  the  exact  data  necessary  are  wanting,  and  in  the 
second  place  the  relationships  in  our  reactions  seem  to  be  satis- 
factorily represented  without  it. 

484.  Valence. — The  idea  just  presented  leads  to  a  simple 
explanation  of  valence  (147,  183).    The  metals  which  form 
only  positive  ions  do  so  by  the  loss  of  one  or  more  electrons  from 
each  atom.     The  valence  of  an  atom  of  a  metallic  element  is 
determined  by  the  number  of  electrons  it  has  lost. 


Electrochemistry 


307 


A  negative  ion,  such  as  Cl  ,  is  an  atom  which  has  taken  up 
an  extra  electron.  Atoms  of  metals  do  not  take  up  extra  elec- 
trons. Only  the  atoms  of  non-metallic  elements  behave  in  this 
way.  The  valence  of  a  negative  ion  consisting  of  one  atom  corre- 
sponds to  the  number  of  electrons  the  atom  has  acquired. 

485.  Theory  of  the  Union  of  Sodium  and  Chlorine. — It  is 
well  known  that  sodium  and  chlorine  unite  very  energetically 
to  form  NaCl.    The  simplest  explanation  of  the  cause  of  union 
is  found  in  the  assumption  that  an  atom  of  sodium  has  a  great 
tendency  to  lose  an  electron,  and  that  an  atom  of  chlorine  has  a 
great  tendency  to  take  up  an  extra  electron.    The  violent  reac- 
tion that  we  observe  when  we  bring  these  two  elements  together 
is  only  the  outward  manifestation  of  the  passage  of  electrons 
from  the  atoms  of  sodium  into  the  atoms  of  chlorine.    The 
residue  of  the  sodium  atom  now  has  an  excess  of  positive  elec- 
tricity, while  the  chlorine  atom  with  its  extra  electron  is  charged 
negatively.     Since  unlike  electric  charges  attract  each  other, 
we  may  well  assume  that  the  two  parts  of  the  NaCl  molecule 
are  held  together  by  electrical  attraction. 

486.  The  Cause  of  lonization. — If  two  insulated  bodies  are 
oppositely  charged,  Fig.  81,  the  force  with  which  they  attract 


FIG.  8 1 


each  other  is  proportional  to  the  product  of  their  charges  and 
inversely  proportional  to  the  square  of  the  distance  between 
them.  There  is,  however,  one  additional  factor  that  determines 
the  strength  of  the  attraction,  and  that  is  the  nature  of  the 
surrounding  medium.  Usually  this  is  air.  If  the  medium  were 


308  Introduction  to  General  Chemistry 

glass  instead  of  air  the  attraction  would  be  only  about  one-third 
as  great,  other  things  remaining  the  same;  but  if  the  medium 
were  water  the  attraction  would  be  only  one-eightieth  as  great 
as  for  air.  If  then  we  dissolve  NaCl  in  water,  the  molecules  are 
surrounded  by  a  medium  which  lessens  enormously  the  attract- 
ive force  which  holds  their  parts  together;  as  the  result,  mole- 
cules will  tend  to  fall  apart,  thus, 

NaCl->Na+  +  Cl-. 

The  positive  part  is  the  sodium  ion,  the  negative  part  the  chlorine 
ion.  According  to  this  explanation  the  molecule  of  salt  before 
it  ionizes  is  made  up  of  two  oppositely  charged  parts.  These 
are  not  ordinary  atoms,  since  the  one  has  lost  an  electron  which 
the  other  has  gained.  We  ought  to  say  that  the  NaCl  mole- 
cule is  made  up  of  a  sodium  ion  (Na+)  electrically  bound  to  an 
ion  of  chlorine  (Cl~).  The  act  of  ionization,  which  takes  place 
when  the  salt  is  dissolved,  is  only  the  falling  apart  of  the  ions  already 
present,  on  account  of  the  great  decrease  in  attractive  force  caused 
by  the  surrounding  water.  In  other  words,  molecules  are  com- 
posed of  bound  ions,  while  in  solution  part  of  the  ions  are  free. 
The  ionization  of  all  acids,  bases,  and  salts  is  explained  in  pre- 
cisely analogous  fashion  to  that  in  the  case  of  NaCl. 

487.  The  Electronic  Description  of  Electrolysis. — According 
to  the  electronic  description  of  electrolysis,  when  an  ion  reaches 
an  electrode  it  either  gains  electrons  or  loses  them.    Thus  the 
positive  ions  Cu++,  Pb++,  H+,  etc.,  each  gain  enough  electrons 
to  make  them  electrically  neutral;   while  Cl~,  I~,  S~~,  and 
O       each  lose  electrons  and  become  free  elements. 

488.  The  Displacement  of  Non-metals  by  One  Another. — It 
will  be  recalled  that  chlorine  acts  on  a  solution  of  hydrobromic 
acid  or  any  bromide,  setting  free  bromine,  thus : 

Cl2+2HBr-»Br2+2HCL  (259) 

Similarly  bromine  acts  on  iodides,  as,  for  example: 
Br2+2KI->I2+2KBr. 


Electrochemistry  309 

Iodine  acts  on  H2S,  in  solution,  setting  free  sulfur: 

(339) 


The  order  in  which  the  four  above-mentioned  elements  displace 
one  another  is  therefore  as  follows: 

Cl,  Br,  I,  S. 

Each  will  set  free  from  its  compounds  any  one  following  it.  We 
may  also  include  fluorine  and  oxygen  in  the  series,  and,  since 
fluorine  will  displace  any  of  the  other  elements  mentioned,  it  will 
head  the  series.  The  position  of  oxygen  is  determined  by  the 
fact  that  a  H2S  solution  reacts  with  atmospheric  oxygen  to  form 
free  sulfur  and  water, 

02+2H2S->2S+2H20,  (339) 

and  that  a  solution  of  HI  also  reacts  with  oxygen  of  the  air 
(slowly)  to  form  water  and  free  iodine, 

02-f4Hfe2l2+2H20.  (265) 

On  the  other  hand,  HBr  solution  is  scarcely  affected  by  oxygen 
gas,  and  HC1  solution  not  at  all.  Oxygen  will  therefore  precede 
iodine  and  sulfur  and  follow  bromine  in  the  list.  The  whole 
displacement  series  is  then  as  follows: 

F,  Cl,  Br,  O,  I,  S. 

489.  Electronic  Interpretation  of  Displacement.  —  If  the  re- 
action 

Cl2+2HBr±sBr2+2HCl 

takes  place  in  very  dilute  solution,  the  two  acids  are  nearly  com- 
pletely ionized,  and  we  may  leave  the  H"1"  ion  out  of  considera- 
tion. The  reaction  in  its  simplest  aspect  is  as  follows: 


This  means  that  each  Br~~  ion  loses  an  electron,  which,  passing 
into  a  chlorine  atom,  changes  the  latter  into  a  Cl"  ion.  We 
conclude  that  chlorine  atoms  take  up  electrons  more  readily  than 


310  Introduction  to  General  Chemistry 

do  atoms  of  bromine.  Considering  next  the  six  elements  of  the 
displacement  series,  we  may  say  that  fluorine  has  the  greatest 
tendency  to  take  up  electrons,  and  sulfur  the  least,  and  that  the 
tendencies  of  th  other  elements  come  in  the  order  indicated  in 
the  list  as  given.  Summarizing,  we  may  say  that  of  two  elements 
of  the  above-mentioned  series  the  one  whose  atoms  have  the  greater 
tendency  to  take  up  electrons  will  set  the  other  free  from  its  compounds 
with  positive  ions. 

490.  The  Displacement  of  Metals  by  One  Another. — Strips 
of  metallic  zinc  placed  in  solutions  of  salts  of  lead,  copper, 
mercury,  and  silver  will  react  as  indicated  by  the  following 

equations: 

Zn+Pb(N03)2->Pb+Zn(NO3)2, 
Zn+  CuSO4->Cu+ZnSO4, 
Zn+HgCl2->Hg+ZnCl2, 

Zn+2AgN03->2Ag+Zn(N03)2. 

In  other  words,  zinc  displaces  each  of  the  above-mentioned  metals 
from  its  salts. 

If  strips  of  metallic  lead  are  placed  in  solutions  of  salts  of 
zinc,  copper,  mercury,  and  silver,  no  reaction  takes  place  with 
the  zinc  salt;  but  the  other  three  metals  are  set  free,  while  the 
lead  atoms  pass  into  solution  as  positive  ions.  In  similar 
fashion  metallic  copper  sets  free  mercury  and  silver  from  their 
salt  solutions,  but  it  does  not  affect  solutions  of  zinc  or  lead  salts. 
Mercury  displaces  silver  from  its  salts  but  has  no  action  on  salts 
of  zinc,  lead,  or  copper.  Metallic  silver  will  not  displace  from 
their  salts  any  of  the  other  metals  just  considered.  The  order 
of  displacement  of  the  five  metals  is  therefore  as  follows : 

Zn,  Pb,  Cu,  Hg,  Ag. 

491.  Electronic  Interpretation  of  Metallic  Displacement. — 

The  action  of  zinc  on  solutions  of  copper  salt.s  may  be  represented 
in  simplified  form  thus: 

Zn+Cu++-»Cu+Zn++. 

This  means  that  an  atom  of  zinc  gives  up  two  electrons  to  an 
atom  of  copper.  Since  zinc  displaces  copper  equally  well  from 


Electrochemistry  311 

solutions  of  all  its  simple  soluble  salts,  we  conclude  that  an 
atom  of  zinc  has  a  greater  tendency  to  lose  electrons  than  has 
an  atom  of  copper,  but,  since  metallic  copper  displaces  silver 
from  any  of  its  salts,  thus, 

Cu+  2  Ag+->2  Ag-h  Cu+ + , 

we  conclude  that  an  atom  of  copper  has  a  greater  tendency  to 
lose  electrons  than  has  an  atom  of  silver. 

The  order  of  the  metals  in  the  displacement  series 

Zn,  Pb,  Cu,  Hg,  Ag 

is  therefore  the  order  in  which  they  fall  according  to  the  decreasing 
ease  with  which  their  atoms  tend  to  lose  electrons.  In  the  case  of 
any  two  metals  of  the  preceding  list  the  one  whose  atoms  have 
the  greater  tendency  to  lose  electrons  will  set  the  other  free  from 
its  compounds  with  negative  ions. 

492.  A  More  Complete  Displacement  Series  of  Metals.— 
Most  of  the  familiar  metals  may  be  included  in  a  single  displace- 
ment series,  which  shows  at  the  same  time  the  tendencies 
of  the  atoms  of  the  metals  to  lose  electrons  and  so  change 
into  positive  ions.  The  list  is  given  in  Table  XIX.  In  this 

TABLE  XIX 

DISPLACEMENT  SERIES  OF  THE  METALS 

Potassium  Nickel 

Sodium  Tin 

Barium  Lead 

Calcium  Hydrogen 

Magnesium  Copper 

Aluminum  Mercury 

Zinc  Silver 

Iron  Platinum 

Cobalt  Gold 

list  (in  which  the  second  column  follows  the  first)  each  metal 
tends  4o  displace ,  or  set  free  from  its  combination  with  negative  ionsy 
any  element  which  follows  it. 

Hydrogen  has  been  placed  in  the  list  between  lead  and  copper. 
Any  metal  above  hydrogen  in  the  series  will  react  with  a  normal 


312  Introduction  to  General  Chemistry 

solution  of  hydrochloric  acid  to  set  free  hydrogen  (at  atmospheric 
pressure)  and  pass  into  solution  as  chloride.  The  metals  follow- 
ing hydrogen  in  the  list  do  not  react  readily,  if  at  all,  with  hydro- 
chloric acid.  The  first  four  elements  of  the  series  react  with 
water  to  set  free  hydrogen.  Therefore  metallic  potassium 
placed  in  a  solution  of  NaCl  does  not  set  free  metallic  sodium 
but  causes  the  evolution  of  hydrogen.  The  order  shown  for 
the  first  four  elements  of  the  table  is  in  fact  that  of  their  tend- 
encies to  lose  electrons  as  determined  by  means  other  than 
direct  displacement. 

493.  The  Production  of  an  Electric  Current.  —  In  the  reac- 
tion between  zinc  and  copper  sulfate  the  essential  change,  as  we 
have  seen,  is  that  represented  by  the  equation 


We  have  said  that  this  change  is  the  result  of  the  passage  of 
two  electrons  from  each  atom  of  zinc  into  an  atom  of  copper. 
Now  if  this  is  true  we  ought  to  be  able  to  get  an  available  elec- 
tric current  from  the  reaction;  but  if  a  piece  of 
zinc  is  dipped  into  a  solution  of  a  copper  salt, 
Fig.  82,  no  evidence  of  the  production  of  such 
a  current  is  to  be  observed.  How  indeed  could 
we  expect  to  detect  the  production  of  an  electric 
current  under  the  conditions  pictured  in  Fig.  82? 
CuSO<  If  a  passage  of  electrons  occurred,  it  would  be 
between  the  zinc  rod  and  the  layer  of  copper 


FIG.  82  i°ns  m  tne  solution  surrounding  the  rod,  and  we 

could  not  readily  detect  such  a  current,  much 
less  make  any  use  of  it.  If  we  wish  to  make  this  supposed 
current  available  for  detection  and  use,  we  must  so  arrange  the 
reacting  substances  that  the  Cu++  ions  are  not  directly  in  con- 
tact with  the  zinc  rod,  and  then  provide  a  wire  for  the  transfer 
of  electrons  from  the  zinc  rod  to  the  copper  ions.  This  can  be 
done  by  arranging  the  four  substances  of  the  reaction 

Zn+  CuS04->ZnSO44-  Cu 


Electrochemistry 


3*3 


in  the  manner  shown  in  Fig.  83 .  Here  we  have  a  zinc  rod  dipping 
into  a  solution  of  ZnSO4  in  one  beaker,  and  a  copper  rod  dipping 
into  a  solution  of  CuSO4  in  the  other  beaker.  A  glass  tube 
filled  with  ZnS04  solution  and  loosely  stoppered  with  cotton 
plugs  forms  a  so-called  salt  bridge  between  the  two  beakers.  If 


FIG.  83 

now  the  two  rods  are  connected  by  wires  to  a  galvanometer,  a 
current  is  found  to  flow  in  a  direction  indicating  the  passage  of 
electrons  from  the  zinc  rod  through  the  wire  (and  galvanometer) 
to  the  copper  rod.  At  the  same  time  metallic  copper  begins  to 
deposit  on  the  copper  plate,  and  metallic  zinc  begins  to  pass  into 
solution.  In  fact,  the  reaction 

Zn-f  CuSO4->ZnSO4+  Cu 

begins  to  take  place  just  as  soon  as  the  metallic  circuit  is  closed 
between  the  upper  ends  of  the  zinc  and  copper  rods.  No  action 
occurs  before  the  circuit  is  closed,  and  all  action  stops  when  the 
circuit  is  broken. 

494.  The  Mechanism,  of  Current  Production. — In  detail  the 
actions  that  occur  with  closed  circuit  are  as  follows :  zinc  atoms 
pass  from  the  rod  into  the  solution,  each  atom  of  zinc  leaving 
behind  two  electrons  and  changing  thereby  into  a  Zn++  ion. 
The  electrons  thus  liberated  flow  through  the  wire  to  the  copper 
rod  in  the  CuS04  solution.  Copper  ions  in  contact  with  the 
copper  rod  take  up  two  electrons  each,  being  thereby  changed 
into  ordinary  copper  atoms.  These  latter  adhere  to  the  copper 
rod  as  a  metallic  coating.  Fresh  Cu++  ions  move  up  to  the 
copper  rod  by  diffusion,  so  that,  as  the  ions  in  contact  with  the 
rod  take  on  electrons  and  change  into  copper  atoms,  others  move 
up  by  reason  of  their  kinetic  motion  to  take  their  places.  On 
the  other  hand  Zn++  ions,  newly  formed  at  the  zinc  rod,  diffuse 


314  Introduction  to  General  Chemistry 

out  into  the  solution.  These  changes  tend  to  cause  a  deficiency 
of  S04~~~  ions  about  the  zinc  rod,  and  an  excess  of  the  same  ions 
about  the  copper  rod.  The  attraction  between  the  excess  of 
SO4  ions,  on  the  one  hand,  and  the  excess  of  Zn+"*~  ions,  on 
the  other,  causes  a  migration  of  these  ions  in  opposite  directions 
through  the  solution  and  the  salt  bridge,  and  thus  serves  to 
maintain  in  each  cubic  centimeter  of  the  whole  solution  as  many 
negative  as  positive  ions  and  thereby  to  keep  the  solution,  as  a 
whole,  electrically  neutral. 

495.  The  Function  of  the  Salt  Bridge. — The  necessity  of 
some  sort  of  connection  between  the  solutions  of  ZnS04  and 
CuS04  in  the  two  beakers,  Fig.  83,  is  obvious.     If  we  remove 
the  salt  bridge,  which  in  this  case  is  a  ZnSO4  solution,  the  circuit 
is  broken,  and  all  action  comes  to  a  stop.     By  the  use  of  the 
bridge  we  are  able,  by  placing  the  CuS04  in  a  separate  beaker, 
to  keep  it  from  coming  in  contact  with  the  Zn  rod.    The  use  of 
a  metal  wire  in  place  of  the  salt  bridge  would  apparently  be  a 
simpler  plan,  but  it  would  not  serve,  because  new  products 
would  be  set  free  by  electrolysis  at  each  end  of  the  wire.    We 
could,  however,  use  in  the  bridge,  instead  of  the  ZnS04,  a  solu- 
tion of  CuS04  or,  in  fact,  of  NaCl  or  almost  any  other  salt.    In 
case  the  bridge  contains  NaCl,  the  Na+  ions  serve  in  place  of 
Zn++   to  carry  the  positive  charge  from  the  ZnS04  solution 
to  the  CuS04  solution,  and  the  Cl"  ions  to  carry  the  negative 
charge  in  the  opposite  direction. 

496.  Galvanic  Cells.    Electric  Batteries. — A  galvanic  cell, 
or,  as  it  is  more  popularly  known,  an  electric  battery,  is  any 
kind  of  apparatus  by  means  of  which  an  electric  current  is 
produced  by  chemical  reactions.    Dry  batteries  and  storage  bat- 
teries are,  at  present,  the  most  familiar  types.    The  first  prac- 
tical form  of  the  zinc-copper  cell  just  studied  was  known  as 
the  Daniell  cell;   a  later  modification  is  known  as  the  gravity 
battery.     Properly  speaking,  the  term  battery  means  a  group 
of  cells,  but  this  term  is  frequently  used  at  present  to  mean  a 
single  cell. 

497.  The    Gravity   Battery. — A   gravity   cell   is    shown   in 
Fig.  84.     It  consists  of  thin  sheets  of  copper  surrounded  by  a 


Electrochemistry  315 

solution  of  copper  sulfate  in  the  lower  half  of  the  glass  jar  and  a 

heavy  zinc  "crowfoot"  surrounded  by  a  zinc  sulfate  solution  in 

the  upper  half.    Attached  to  the  copper  sheets  is  an  insulated 

copper  wire.    A  new  cell  is  set  up  by  rilling  the  jar  with  water, 

placing  the  copper  and  zinc  in  position,  and  adding  more  than 

enough  solid  CuS04  (blue  vitriol  or  bluestone)  to  saturate  the 

lower  layer.    No  ZnS04  need  be  added;  instead, 

twenty  or  thirty  g.  of  NaCl  are  sprinkled  into 

the  water.    The  solution  is  not  stirred.    The 

CuS04  gradually  dissolves,  giving  a  saturated 

solution  which  soon  fills  the  lower  part  of  the 

cell.     If  now  the  insulated  wire  leading  from 

the  copper  is  connected  to  the  zinc,  a  current 

flows  through  the  wire,  and  the  changes  already 

described   take  place.     The  NaCl  is  used  to  FIG  g 

make    the    water    conduct    the   current  prior 

to  the  formation  of  sufficient  ZnS04  for  this  purpose.     Until 

recently  the  gravity  battery  was  used  to  operate  all  telegraph 

lines. 

498.  Other  Kinds  of  Galvanic  Cells. — It  is  possible  to  make 
a  cell  that  will  give  a  current  by  the  use  of  any  pair  of  metals 
(not  acted  upon  by  water) ,  each  surrounded  by  a  solution  of  one 
of  its  salts.     In  each  cell  the  experimental  arrangement  may  be 
that  shown  in  Fig.  83. 

499.  Electromotive  Force  and  Voltage. — A  body  at  rest  can 
be  set  in  motion  only  by  the  action  of  a  force  (Newton's  law). 
In  a  similar  manner  we  assume  that  the  current  (stream  of 
electrons)  produced  by  a  battery  is  the  result  of  an  electrical 
force   called   the    electromotive  force,   E.M.F.    The   unit  of 
E.M.F.  is  the  volt  (named  after  the  pioneer  electrical  experi- 
menter Volta).    The  gravity  cell  has  an  E.M.F.  of  i .  i  volts. 

The  farther  the  two  metals  forming  the  electrodes  of  a  cell 
of  any  kind  are  removed  from  each  other  in  the  displacement 
series  (492)  the  greater  the  E.M.F.  of  the  cell.  The  reason  for 
this  is  found  in  the  fact  that  the  metals  heading  the  list  give 
off  electrons  most  readily  (with  greatest  force).  The  order  in 
the  list  represents,  in  fact,  the  relative  force  with  which  the 


316  Introduction  to  General  Chemistry 

element  loses  electrons.  The  difference  of  such  forces  for  the 
two  metals  of  a  cell  is,  for  practical  purposes,  the  chief  deter- 
mining factor  of  the  E.M.F.  (voltage)  of  the  cell.  This  difference 
of  forces  between  the  electrodes  is  also  often  called  the  potential 
difference  of  the  electrodes. 

There  is  another  important  factor  to  be  considered  besides 
the  nature  of  the  reactions  at  the  electrodes,  and  that  is  the  con- 
centration of  the  ions  in  solution.  For  example,  the  more  con- 
centrated the  copper  ions  at  the  copper  electrode  the  faster  is  the 
reaction  carried  on  by  these  ions  at  a  given  temperature.  Now 
the  difference  of  potential  at  the  terminals  of  a  cell  is  a  measure 
of  the  rate  of  the  reaction  in  progress;  hence  it  will  be  increased 
or  decreased  by  concentration  changes  in  the  solutions.  To 
make  careful  comparison  of  the  electromotive  forces  of  cells  the 
concentrations  of  the  ions  must  therefore  be  taken  strictly  into 
account.  However,  in  the  series  we  are  considering,  no  moderate 
variation  of  the  concentrations  of  the  ions  from  those  found  in 
the  ordinary  laboratory  reagents  (o.oi  to  6N  approximately) 
will  produce  results  different  from  those  described  here,  in  the 
cases  under  consideration.  The  effect  of  the  concentration  of 
ions  on  cell  potentials  should  be  considered  in  an  exact  study  of 
the  latter  subject. 

500.  Electrical  Energy. — Electrical  energy  always  depends 
on  two  factors,  voltage  and  quantity  of  electricity.  The  unit 
of  electrical  energy  is  the  joule,  named  after  J.  P.  Joule,  the 
celebrated  English  scientist,  whose  work  on  the  mechanical 
equivalent  of  heat  was  discussed  earlier  (370).  One  joule  is  the 
amount  of  energy  produced  when  a  quantity  of  one  coulomb  of  elec- 
tricity flows  through  a  conductor,  the  ends  of  which  have  a  potential 
difference  of  one  wit.  In  general,  joules  =  voltsX  coulombs. 
For  example,  if  a  gravity  cell  of  i .  i  volts  E.M.F.  delivers  10 
coulombs,  the  electrical  energy  produced  is  i .  iX  10  =  n  joules. 
Since  the  joule  is  an  energy  unit,  its  value  is  expressible  in  other 
energy  units.  Careful  experiment  has  shown  that 

i  joule  =  10,200  g.cm., 
i  joule = o .  24  calorie, 
i  calorie =4 . 18  joules. 


Electrochemistry 


It  is  electrical  energy  which  a  consumer  pays  for  and  uses. 
The  same  number  of  electrons  go  back  to  the  positive  pole  of  a 
battery  as  leave  the  negative  pole,  but  they  lose  energy  in  so 
going.  The  energy  which  the  electrons  give  up  may  be  liberated 
as  heat  or  may  be  converted  into  work  by  means  of  devices  like 
the  motor. 

501.  Electronic  Explanation  of  Oxidation-Reduction  Reac- 
tions. —  The  action  of  chlorine  on  ferrous  chloride  in  solution 
(i?3>  332)  is  a  simple,  typical  example  of  an  oxidation-reduction 
reaction, 

2FeCl2-fCl2->2FeCl3. 

This  reaction  in  dilute  solution  may  be  represented  by  the  simpli- 
fied equation 


The  ferrous  ion,  Fe++,  which  is  the  reducing  agent,  is  oxidized 
to  Fe+++  by  the  chlorine  atom,  which  is  the  oxidizing  agent. 
This  is  explained  by  assuming  that  the  Fe++  ion  (which  is  an 
iron  atom  that  has  already  lost  two  electrons)  gives  up  a  third 
electron,  which,  passing  into  the  Cl  atom,  changes  the  latter 
into  a  Cl~  ion.  Thus  we  see  that  the  oxidation  of  the  Fe++  ion 


FIG.  85 

consists  in  its  loss  of  an  electron;  and  the  reduction  of  the  Cl  atom 
consists  in  its  gain  of  an  electron. 

502.  Oxidation-Reduction  Cells. — The  transfer  of  electrons 
which  occurs  in  the  reaction  just  studied  can  be  made  to  yield 
an  available  electric  current  quite  as  readily  as  that  which 
takes  place  in  the  reaction  between  metallic  zinc  and  copper 
sulfate  (493).  We  may  demonstrate  this  fact  by  means  of  the 
arrangement  shown  in  Fig.  85.  Platinum  electrodes  are  placed 


318  Introduction  to  General  Chemistry 

in  each  of  two  beakers,  one  of  which  contains  the  FeCl2  solution, 
the  other  the  C12  solution  (together  with  some  FeCl3  or  NaCl  to 
make  the  solution  conduct).  A  salt  bridge  joins  the  two  solu- 
tions. Wires  from  the  electrodes  are  connected  with  a  galvano- 
nometer,  which  shows  the  passage  of  a  current  in  a  direction 
indicating  a  flow  of  electrons  in  the  wire  from  the  electrode  in  the 
FeCl2  solution  to  that  in  the  C12  solution.  The  platinum  elec- 
trodes serve  as  carriers  of  electrons  into  and  out  of  the  solutions. 
Platinum  is  superior  to  any  other  metal  except  gold  for  this 
purpose,  because  of  its  very  slight  tendency  to  pass  into  solution 
as  ions. 

503.  Further  Examples  of  Oxidation-Reduction  Reactions.  — 
Oxidation-reduction  reactions  are  very  common.  They  may 
all  be  interpreted  in  terms  of  electron  transfers,  as  the  following 
additional  examples  will  illustrate.  Ferric  sulfate  is  reduced  by 
zinc  according  to  the  equation 

Fe2(S04)3+Zn-»2FeSO4+ZnSO4.  (335) 

The  simplified  ionic  equation  is 


Each  atom  of  zinc  loses  two  electrons  and  changes  into  a  Zn++ 
ion;  these  two  electrons  are  taken  up,  one  by  each  Fe+++  ion, 
which  is  thereby  changed  to  a  Fe++  ion.  The  zinc,  which  loses 
electrons,  is  the  reducing  agent  and  is  oxidized  by  ferric  ions, 
which  gain  electrons  and  are  thereby  reduced  to  ferrous  ions. 

The  action  of  ferric  salts  and  soluble  iodides  is  illustrated  by 
the  following  reaction: 

2FeCl3+2KI^2FeCl2+2KCl+I2, 
or  in  simplified  form  by 


A  closely  analogous  reaction  takes  place  in  the  reduction  of 
ferric  salts  by  hydrogen  sulfide: 

2FeCL+HaS->2FeCla+2HCl+S. 


Electrochemistry  319 

The  simplified  equation  is 


The  electronic  explanation  of  each  of  the  foregoing  reactions  can 
easily  be  made  by  the  student. 

Other  more  intricate  oxidation-reduction  reactions,  which 
will  require  a  somewhat  more  extended  discussion,  will  be  taken 
up  in  subsequent  chapters. 

In  all  oxidation-reduction  reactions  transfers  of  electrons 
occur;  and  in  all  cases  the  atom  or  ion  which  is  oxidized  loses 
one  or  more  electrons,  and  the  atom  or  ion  which  is  reduced  gains 
one  or  more  electrons.  //  an  ion  does  not  change  its  charge  or 
its  composition  in  the  course  of  a  reaction  it  is  neither  oxidized  nor 
reduced. 

504.  The   Oxidation  and  Reduction  of  Metals.  —  When  a 
metal  passes  into  solution  its  atoms  take  on  positive  charges. 
This  means  that  each  atom  of  a  metal  loses  one  or  more  elec- 
trons when  it  changes  into  an  ion.    Since  we  have  defined 
oxidation  as  the  loss  of  electrons  (501)  we  can  therefore  say 
that  when  a  metal  changes  into  its  ions  it  is  oxidized.     For 
example,  in  the  reaction 

Fe+  CuS04->FeS04+  Cu, 

which  we  may  write  in  simplified  form  thus, 

Fe+Cu++->Fe++-fCu, 

we  say  that  the  metallic  iron  is  oxidized  to  ferrous  ions,  and  the 
copper  ions  are  reduced  to  metallic  copper.  We  have  already 
seen  that  the  further  oxidation  of  Fe++  to  Fe+++  involves  a 
loss  of  one  additional  electron. 

505.  The  Oxidation  and  Reduction  of  Non-metals.—  When 
a  non-metal  (chlorine  for  example)  passes  into  solution,  its  atoms 
take  on  electrons.    We  say,  therefore,  that  in  such  a  case  the 
element  is  reduced.     Conversely  we  say  that  its  ions  are  oxidized 
when  by  loss  of  electrons  they  are  changed  to  atoms  of  the 
element. 


320 


Introduction  to  General  Chemistry 


506.  Oxidation-Reduction     Potentials. — Every     oxidation- 
reduction  reaction  can  by  suitable  arrangement  be  made  to 
yield  an  electric  current.     The  E.M.F.  (voltage)  of  an  oxidation- 
reduction  cell  is  the  measure  of  the  force  with  which  the  reaction  tends 
to  take  place.    The  stronger  the  oxidizing  tendency  of  the  oxidiz- 
ing agent  and  the  stronger  the  reducing  tendency  of  the  reducing 
reagent  the  greater  the  E.M.F.  of  the  cell.    A  systematic  study 
of  such  cells  has  shown  that  all  oxidizing  and  reducing  agents  may 
be  arranged  in  a  series  in  the  order  of  their  decreasing  oxidizing 
tendencies  and  increasing  reducing  tendencies. 

507.  Oxidation  and  Reduction  by  Means  of  the  Electric 
Current. — We  have  shown  that  oxidation  and  reduction  are 
capable  of  producing  electric  currents.    There  now  remains 
to  show  that  an  electric  current  can  accomplish  oxidation  and 
reduction.     Two   beakers,   Fig.   86,   are  fitted   with  platinum 


Fed, 


FIG.  86 

electrodes  and  joined  with  a  salt  bridge,  and  in  one  is  placed  a 
solution  of  FeCl3,  and  in  the  other  HCI.  Upon  passing  a  current 
from  a  battery  of  two  dry  cells  so  connected  that  the  electrode 
in  the  FeCl3  will  be  the  cathode,  it  will  be  found  that  the  FeCl3 
is  reduced  to  FeCl2,  while  at  the  same  time  HCI  is  oxidized  to 
free  chlorine  at  the  anode.  The  explanation  is  as  follows:  The 
battery  sends  a  steady  stream  of  electrons  through  the  wire  to 
the  cathode;  one  electron  passes  from  the  latter  into  each 
Fe+++  ion  coming  in  contact  with  it,  changing  the  Fe+++ 
into  Fe++.  At  the  anode  Cl~  ions  coming  in  contact  with 
this  electrode  give  up  to  it  their  electrons  and  change  thereby 
into  ordinary  Cl  atoms.  The  latter  then  unite  in  pairs  to  form 
molecules,  aggregates  of  which  soon  form  bubbles  that  escape 
into  the  air.  In  the  solution  Fe+++  ions  are  attracted  by  and 
migrate  toward  the  cathode,  while  Cl"  ions  are  attracted  by  and 


Electrochemistry  321 

migrate  toward  the  anode.  Thus  the  transfer  of  electricity  from 
one  electrode  to  the  other  in  the  solution  is  accomplished  by  means 
of  the  moving  ions,  while  in  the  wire  we  have  a  stream  of  electrons 
set  in  motion  by  the  battery.  A  great  variety  of  other  oxidations 
and  reductions  in  solution  can  be  accomplished  by  means  of  the 
electric  current.  In  fact,  since  we  may  consider  the  change  of 
a  metal  into  its  ions  as  an  oxidation  of  the  former,  and  the  reverse 
change  a  reduction  of  the  ions,  we  may  go  farther  and  say  that 
all  processes  of  electrolysis  result  in  oxidation  and  reduction.  The 
anode  is  the  seat  of  oxidation,  since  it  takes  up  electrons;  the 
cathode  is  the  seat  of  reduction,  because  it  furnishes  electrons. 
These  statements  apply  to  all  electrolyses  irrespective  of  whether 
the  substances  formed  or  liberated  at  the  electrodes  are  elements 
or  compounds. 

508.  The  Conversion  of  Chemical  Energy  into  Electrical 
Energy. — The  production  of  heat  by  a  chemical  reaction  has 
been  explained  (373)  as  being  due  to  the  conversion  of  chemical 
energy  of  the  reacting  substances  into  heat  energy.  If  metallic 
zinc  is  placed  in  a  solution  of  copper  sulfate  so  that  the  reaction 

Zn+  CuSO4-»ZnSO4+  Cu 

takes  place  without  the  production  of  an  available  electric 
current,  the  quantity  of  heat  liberated  is  50,100  calories  for  one 
symbol  weight  of  zinc.  If  the  same  amount  of  zinc  reacts  with 
copper  sulfate  in  a  gravity  cell,  2X96,500  coulombs  of  elec- 
tricity are  delivered  into  the  circuit  at  an  E.M.F.  of  1.09  volts. 
The  electrical  energy  produced  is  2X96,500X1.09  =  210,400 
joules.  Since  i  calorie  =  4.18  joules,  210,400  joules  =  2 10,400 
-7-4.18  =  50,300  calories.  Thus  we  see  that  electrical  energy 
equivalent  to  50,300  calories  is  produced  in  a  cell,  instead  of 
50,100  calories  of  heat  produced  when  the  same  amounts  of  the 
substances  react  directly,  without  the  production  of  a  current. 
The  small  excess  of  energy  produced  in  a  cell  is  accounted  for  by 
the  fact,  established  by  experiment,  that  this  amount  of  energy 
is  absorbed  as  heat  from  the  surroundings  as  the  cell  operates. 
Similar  results  are  observed  in  the  energy  production  of  all  other 
galvanic  cells.  The  electrical  energy  produced  by  any  cell  is 


322  Introduction  to  General  Chemistry 

equal  to  the  chemical  energy  liberated  or  lost,  plus  or  minus  an 
additional  amount  of  energy — plus  if  heat  is  taken  up  from  the 
surroundings  and  minus  if  it  is  given  out  to  the  surroundings.  We 
may  consider  a  galvanic  cell  or  battery  merely  as  a  device  for 
converting  chemical  energy  into  electrical  energy. 

509.  Conversion  of  Electrical  Energy  into  Chemical  Energy. — 
In  all  processes  of  electrolysis  electrical  energy  is  used  up  in  the 
production  of  new  chemical  substances,  and  we  may  conclude 
at  once  that  the  electrical  energy  used  is  changed  to  chemical 
energy. 


CHAPTER  XXI 
NITROGEN  AND  AMMONIA 

510.  Introduction. — Many  facts  about  nitrogen  and  two  of 
its  important  compounds,  ammonia  and  nitric  acid,  have  already 
been  discussed.  We  may  recall  that  about  four-fifths  by  volume 
of  air  is  uncombined  nitrogen,  which  is  left  in  a  nearly  pure  state 
as  a  colorless,  odorless  gas  when  air  is  freed  from  oxygen  (15). 
Pure  nitrogen  is  obtained  by  passing  ammonia  over  red-hot 
copper  oxide  (84) : 

2NH3+3CuO->3Cu+3H20+N2. 

The  symbol  weight  of  nitrogen  is  14,  and  since  22.4  liters  of  the 
gas  at  o°  and  76  cm.  weigh  28  g.,  the  formula  is  N2  (75).  This 
means  that  a  molecule  of  nitrogen  consists  of  two  atoms  (215). 

Ammonia,  NH3,  is  a  colorless  gas  which  can  be  made  by  the 
union  of  nitrogen  and  hydrogen,  by  the  action  of  electric  sparks, 

N2+3H2^2NH3.  (298) 

It  has  a  very  penetrating  odor  and  dissolves  easily  in  water, 
forming  at  the  same  time  some  ammonium  hydroxide: 

NH3+H2O->NH4OH.  (91) 

Ammonia  and  ammonium  hydroxide  may  be  completely  elimi- 
nated from  a  solution  by  boiling  the  latter. 

Ammonium  hydroxide  is  a  base  which  neutralizes  acids  to 
form  salts,  among  which  we  have  studied  the  chloride  NH4C1 
(92),  the  nitrate  NH4N03  (105),  the  sulfate  (NH4)2SO4  (101), 
and  the  acid  sulfate  NH4HSO4  (101). 

Ammonium  hydroxide  is  a  very  much  weaker  base  than 
sodium  hydroxide  (409,  429);  the  latter  substance  reacts  with 
ammonium  salts  to  set  free  NH3,  as  illustrated  by  the  following 

equation: 

NH4Cl+NaOH->NaCl+H2O-hNH3.  (426) 

323 


324  Introduction  to  General  Chemistry 

511.  The  Occurrence  of  Nitrogen. — Besides  the  occurrence 
of  free  nitrogen  in  air,  this  element  is  also  found  in  nature  as  a 
constituent  of  many  compounds,  among  which  are  the  familiar 
substances  ammonia  and  the  nitrates  of  sodium  and  potassium. 
Nitrogen  is  also  an  essential  constituent  of  the  proteins,  which 
constitute  the  bulk  of  all  kinds  of  flesh.     Proteins  are  also  present 
in  plants,  particularly  in  their  seeds.     Cereals  are  especially 
rich  in  nitrogenous  matter.     In  the  course  of  the  decay  of  animal 
and  vegetable  matter  the  nitrogen  is  changed  first  into  NH3  and 
finally  into  nitrates. 

512.  The  Element  Nitrogen. — Nitrogen  was  discovered  by 
Rutherford,  professor   of    chemistry   in    Edinburgh,  in    1772. 
Lavoisier,  who  was  the  first  to  recognize  it  as  an  element,  called 
it  azote.    Its  English  name,  nitrogen,  suggested  by  Chaptal, 
indicates  that  the  element  can  be  made  from  saltpeter,  KNO3) 
for  which  the  Greek  name  was  nitron. 

513.  The  Preparation  of  Nitrogen. — So-called  atmospheric 
nitrogen,  obtained  by  removing  oxygen,  carbon  dioxide,  and 
water  from  air,  is  not  pure  and  cannot  easily  be  freed  from  its 
residual  impurities,  amounting  to  about  i  per  cent,  which  con- 
sist of  chemically  inactive  gases,  chiefly  argon.    Atmospheric 
nitrogen  is  prepared  on  a  technical  scale  by  the  Linde  process, 
in  which  air  is  first  liquefied  and  the  more  volatile  nitrogen 
distilled  from  the  less  volatile  liquid  oxygen.    Pure  nitrogen 
can  be  made  not  only  from  NH3  and  CuO  but  also  in  several 
other  ways.    One  of  the  best  of  these  is  the  following:    20  g. 
of  NH4C1  and  25  g.  of  sodium  nitrite,  a  salt  having  the  formula 
NaN02,  and  50  c.c.  of  water  are  placed  in  a  flask  (Fig.  87)  and 
heated  gently,  whereupon  the  following  reaction  occurs: 

NH4Cl+NaNO2->NaCl+N2+  2H2O. 

This  reaction  takes  place  in  two  stages,  of  which  the  first  is 
NH4Cl+NaNO2->NaCl+NH4NO2. 

The  salt  NH4NO2,  ammonium  nitrite,  is  unstable  and  decomposes 
at  once  into  N2  and  H2O,  thus: 

NH4NO2->N2+2H2O. 


Nitrogen  and  Ammonia 


325 


514.  The  Properties  of  Nitrogen.— Nitrogen  is  a  colorless, 
odorless  gas,  as  might  be  inferred  from  the  fact  that  four-fifths 
by  bulk  of  the  air  is  free  nitrogen.    The  gas  does  not  support 
combustion,  nor  does  it  burn  in  the  ordinary  sense  of  the  term. 
However,  nitrogen  unites  with  oxygen,  under  certain  conditions, 
to  form  the  oxides  NO  and  N02.     These  very  important  reactions 
will  be  considered  in  detail  in  chapter  xxii,  as  by  their  means 
nitric  acid  can  be  made  from  the  nitrogen  of  the  air. 

Nitrogen  unites  di- 
rectly with  several  ele- 
ments, mostly  metals, 
and  especially  at  high 
temperatures,  to  form 
nitrides,  among  which 
calcium  nitride,  Ca3N2; 
magnesium  nitride, 
Mg3N2;  and  aluminum  FIG.  87 

nitride,  A1N,  may  be 

mentioned.  Nitrides  usually  react  readily  with  water  to  form 
ammonia  and  the  hydroxide  of  the  metal: 

A1N+3H2O->A1(OH)3+NH3. 

A  reaction  of  great  technical  importance  takes  place  when 
nitrogen  gas  is  passed  over  calcium  carbide,  CaC2  (49),  at  a 

white  heat: 

CaC2+N2->CaCN2+C. 

The  product,  CaCN2,  is  calcium  cyanamid;  it  is  also  called 
nitrolime.  It  is  a  valuable  fertilizer  and  is  also  used  in  the  manu- 
facture of  ammonia  and  cyanides  by  methods  which  will  be 
described  later. 

515.  The  Assimilation  of  Nitrogen  by  Plants. — All  plants 
require  for  their  growth  nitrogen  in  some  form,  in  addition  to 
other  things.    Fertile  soils  contain  considerable  nitrogen  in  the 
form  of  compounds,  among  which  are  ammonium  salts,  nitrates, 
and  decaying  vegetable  and  animal  matter.     Such  compounds 
supply  plants  with  their  required  nitrogen,  since  in  general  plants 
can  assimilate  only  combined  and  not  free  nitrogen.     A  few 


326  Introduction  to  General  Chemistry 

plants,  however,  among  which  are  clover,  alfalfa,  and  other 
legumes,  have  the  power  of  utilizing  uncombined  nitrogen  from 
air  which  has  permeated  the  soil.  This  they  do  through  the 
medium  of  root  nodules,  which  are  bacterial  growths,  whitish 
in  color  and  often  as  large  as  peas.  Free  nitrogen  is  taken  up 
by  the  bacteria  in  the  nodules  and  changed  into  compounds, 
which  then  serve  as  a  source  of  nitrogen  for  the  plant.  A  large 
part  of  the  nitrogen  taken  up  by  plants  is  concentrated  in  their 
seeds  as  complex  compounds  called  proteins.  Wheat,  for 
example,  contains  10  to  15  per  cent  of  proteins. 

516.  The  Sources  of  Ammonia. — Although  ammonia  occurs 
in  very  minute  amounts  in  the  air,  in  natural  waters,  and  in  the 
soil,  being  formed  by  the  decay  of  animal  and  vegetable  matter, 
such  occurrences  are  not  a  practical  source  of  supply.     In  earlier 
times  ammonia,  then  called  spirits  of  hartshorn,  was  made  by 
heating  bones,  hoofs,  horns,  and  other  animal  matter  with  lime 
(calcium  oxide).    At  present  ammonia  is  produced  in  large 
amount  as  a  by-product  in  the  distillation  of  coal  for  the  pro- 
duction of  fuel  and  illuminating  gas  and  coke.     Coal  usually 
contains  i  or  2  per  cent  of  combined  nitrogen,  and  when  heated 
to  a  high  temperature  with  exclusion  of  air,  the  nitrogen  com- 
pounds present  are  decomposed,  with  the  formation  of  ammonia, 
which  passes  off  with  the  gas.    The  ammonia  is  absorbed  in 
dilute  sulfuric  acid,  giving  a  solution  which  upon  evaporation 
yields  crystals  of  ammonium  sulf ate ;  and  this  salt,  when  heated 
with  lime,  yields  free  ammonia: 

(NH4)2SO4+  CaO^CaSO4+ H2O+  2NH3 . 

517.  The  Physical  Properties  of  Ammonia. — For  the  sake 
of  ready  reference  we  may  summarize  here  the  more  important 
data  regarding  ammonia.    It  is  a  colorless  gas  of  penetrating 
odor;  one  liter  at  o°  and  76  cm.  weighs  o .  772  g. ;  and  22.4  liters 
weigh  17  grams.    Upon  being  strongly  compressed  ammonia 
condenses  to  a  colorless  liquid.     Liquid  ammonia  is  an  important 
article  of  commerce;  it  is  marketed  in  steel  cylinders  of  about 
two  hundred  pounds  capacity.    At  20°  the  liquid  has  a  vapor 
pressure  of  8.4  atmospheres.     If  this  pressure  is  released  the 


Nitrogen  and  Ammonia  327 

liquid  boils;  that  is,  it  changes  into  gaseous  NH3,  and  its  tem- 
perature falls  to  —33°  (519),  which  is  its  boiling-point  under  one 
atmosphere  pressure. 

At  20°  water  dissolves  about  seven  hundred  times  its  volume 
of  NH3  gas,  part  of  it  going  to  form  the  hydroxide  NH4OH, 
which  is  a  weak  or  slightly  ionized  base.  Pure  ammonium 
hydroxide  of  commerce  has  a  specific  gravity  of  0.90  and  con- 
tains 28  per  cent  by  weight  of  NH3. 

518.  The  Uses  of  Ammonia. — Ammonia  is  a  substance  of 
fundamental  importance.     It  is  made  and  consumed  in  immense 
quantities,  and  its  uses  are  numerous.     In  the  household  it  is 
added  to  water  used  in  washing.     In  laboratories  and  chemical 
works  it  is  indispensable.    Ammonium  salts,  particularly  the 
sulfate,  are  used  in  enormous  quantities  in  fertilizers.     The 
manufacture  of  nitric  acid  from  ammonia  has  become  since  the 
beginning  of  the  war  a  matter  of  vital  importance;  this  subject 
will  be  taken  up  in  the  next  chapter  (570).     The  use  of  ammonia 
in  making  ice  may  now  be  considered  in  some  detail,  as  it  is  not 
only  of  commercial  importance  but  also  of  much  scientific 
interest. 

519.  The  Manufacture  of  Artificial  Ice. — It  will  be  recalled 
that  trie  change  of  water  into  steam  at  100°  takes  place  with 
the  absorption  of  540  calories  of  heat  per  gram   (115);    and 
further  that  "every  pure  liquid  has  a  latent  heat  of  evaporation." 
The  latent  heat  of  evaporation  of  liquid  ammonia  at  its  boiling- 
point,    —  33°,  is  258  calories.     This  means  that  liquid  NH3 
released  from  pressure  passes  into  gas  with  an  absorption  of  258 
calories  of  heat  per  gram.    This  heat  is  taken  from  the  ammonia 
and  the  vessel  containing  it,  and  if  this  vessel  is  immersed  in 
water  the  latter  is  quickly  frozen.     Since  the  conversion  of  i  g. 
of  water  at  o°  into  ice  requires  an  absorption  of  79  calories, 
enough  heat  is  absorbed  by  the  evaporation  of  i  g.  of  NH3  to 
freeze  about  3  g.  of  water. 

The  practical  application  of  these  principles  to  the  manu- 
facture of  artificial  ice  is  illustrated  diagrammatically  by  Fig.  88. 

The  pump,  A,  draws  ammonia  gas  from  the  pipes  immersed 
in  the  tank,  0,  and  compresses  the  gas  sufficiently  so  that  when 


328 


Introduction  to  General  Chemistry 


the  latter  is  cooled  by  running  water  in  the  condenser,  B, 
liquid  ammonia  is  formed,  and  collects  in  receiver,  C.  From 
C  liquid  ammonia  flows  into  the  pipes  that  run  through  the 
tank,  O.  In  the  pipes  the  ammonia  evaporates,  thus  cooling 
the  brine  (freezing-point  —20°)  with  which  0  is  filled.  The 
water  to  be  frozen  is  contained  in  cans  of  200  pounds  or  more 
capacity,  immersed  in  the  brine. 


FIG.  88 


520.  The  Synthesis  of  Ammonia. — We  have  already  learned 
that  NH3  can  be  made  from  nitrogen  and  hydrogen  by  passing 
an  electric  spark  through  the  mixed  gases  (298).  The  electric 
spark  is  effective  mainly  as  a  source  of  heat: 

3H2+N2->2NH3. 

This  reaction  reaches  a  state  of  equilibrium  when  but  very  little 
NH3  has  been  formed,  for  the  reason  that  the  reverse  action, 
the  decomposition  of  NH3  into  H2  and  N2,  takes  place  very 
easily.  At  a  temperature  of  800°  equilibrium  is  reached  when 
only  one  molecule  of  NH3  remains  undecomposed  out  of  every 
10,000  taken.  At  first  thought  it  would  seem  to  be  quite 
hopeless  to  try  to  make  use  of  this  reaction  as  a  means  of  manu- 
facturing NH3  from  H2  and  N2,  because  of  the  very  minute  pro- 
portion of  these  gases  that  unite  before  the  reaction  reaches 
equilibrium.  In  order  that  a  process  based  on  this  reaction 
shall  be  a  practical  success,  substantially  all  of  the  H2  and  N2 
taken  must  be  converted  into  NH3.  The  big  problem  then  is 
to  obtain  from  the  H2  and  N2  taken  the  other  9,999  molecules 
of  NH3.  On  account  of  the  great  technical  importance  of  the 
subject  and  as  a  beautiful  illustration  of  the  way  in  which 


Nitrogen  and  Ammonia  329 

fundamental  chemical  principles  (laws)  are  used  in  working 
toward  desired  results,  we  shall  consider  this  problem  in  detail. 
521.  The  Effect  of  Temperature  on  the  Equilibrium. — We 
have  already  learned  (288)  that  for  any  reaction  the  proportions 
of  the  substances  present  when  equilibrium  is  reached  are 
different  at  different  temperatures.  With  change  of  tempera- 
ture the  equilibrium  of  any  reaction  is  shifted  in  a  way  that  can 
be  definitely  predicted  from  its  heat  of  reaction.  If  heat  is 
given  out  when  a  substance  is  formed,  heat  will  be  absorbed 
when  the  same  substance  is  decomposed.  When  we  raise  the 
temperature  of  a  system  in  equilibrium,  the  state  of  equilibrium 
shifts  in  such  a  way  that  heat  is  absorbed  (288,  367).  The  heat 
of  formation  (361)  of  NH3  is  12,200  calories: 

3H2+N2->2NH3+2Xi2,2oo  cal. 

Therefore  the  higher  the  temperature  the  smaller  the  propor- 
tion of  NH3  present  in  the  equilibrium  mixture,  and  vice  versa, 
as  shown  by  the  following  table,  in  which  is  given  the  percentage 
of  NH3  at  various  temperatures  and  one  atmosphere  pressure: 

Temperature  Per  Cent  of  NH3 

800° o.oi 

600° 0.05 

400°   .......   0.48 

It  is  evident  that  the  lower  the  temperature  the  more  favorable 
the  result;  but  a  serious  practical  difficulty  is  encountered  if 
one  tries  to  make  NH3  at  low  temperatures;  the  union  of  H2 
and  N2  proceeds  the  more  slowly  the  lower  the  temperature,  so 
that  at  400°  it  would  require  days  for  equilibrium  to  be  reached. 
In  this  respect  this  reaction  is  like  all  others;  the  speed  of  reac- 
tion, other  things  being  equal,  is  slower  the  lower  the  tempera- 
ture. The  experimenter  then  faces  this  dilemma:  at  a  high 
temperature  the  reaction  gives  very  little  NH3;  at  a  low  tem- 
perature it  goes  too  slowly.  Is  there  no  way  to  surmount  this 
difficulty?  Is  there  no  way  to  hasten  a  reaction  except  by 
increasing  the  temperature?  Yes,  there  is,  and  this  by  means 
of  a  catalytic  agent ! 


330  Introduction  to  General  Chemistry 

522.  Catalytic  Agents  for  the  Ammonia  Reaction. — As  illus- 
trations of  cases  of  catalysis  we  may  recall  that  CuCl2  catalyzes 

the  reaction 

4HC1+O2-»2H2O+  2C12 ;  (239) 

that  Mn02  promotes  the  decomposition  of  KC1O3  (306);  and 
that  platinum  black  causes  H2  and  02  to  unite  rapidly  at  room 
temperature  (303).  Therefore  it  will  not  be  surprising  to  learn 
that  several  catalytic  agents  were  found  which  hasten  the 
ammonia  reaction  enormously;  among  such  were  the  metals 
iron,  manganese,  and  uranium.  By  the  use  of  these  and  prob- 
ably other  substances,  the  nature  of  which  has  not  yet  been 
disclosed  by  the  discoverers,  the  ammonia  reaction  is  caused  to 
reach  equilibrium  rapidly  at  400°  to  500°;  but  even  at  400° 
only  one-half  of  i  per  cent  of  ammonia  is  formed.  In  the  next 
paragraph  we  shall  see  how  the  equilibrium  can  by  another  means 
be  shifted  still  farther  in  the  desired  direction. 

TABLE  XX 

PERCENTAGES  OF  AMMONIA  AT  VARIOUS 
TEMPERATURES  AND  PRESSURES 


Temperature 

Percentage  at 
i  Atmosphere 

Percentage  at 
100  Atmospheres 

800°  

O.OI 

I  .  I 

Itt0'.'. 
500°  

O.O2 
0.05 
0.13 

2.  I 

4-5 
10.8 

523.  The  Effect  of  Pressure  on  the  Equilibrium. — In  chapter 
xiii  (287)  we  learned  that  "  the  effect  of  increase  of  pressure  on 
any  system  in  equilibrium  is  in  all  cases  to  shift  the  equilibrium 
so  as  to  favor  the  formation  of  substances  occupying  a  smaller 
volume."  In  the  reaction 

3H2+N2->2NH3 

three  volumes  of  hydrogen  and  one  volume  of  nitrogen,  if  com- 
pletely united,  will  yield  two  volumes  of  ammonia  (77),  from 
which  we  conclude  that  an  increase  of  pressure  will  shift  the 
equilibrium  so  as  to  favor  the  formation  of  more  NH3.  Table 
XX  gives  the  results. 


Nitrogen  and  Ammonia  331 

It  has  been  found  practicable  to  work  at  pressures  of  100 
atmospheres  (1,400  Ib.  per  square  inch)  and  even  higher,  and 
thus  greatly  to  increase  the  proportion  of  NH3  formed.  It  now 
remains  to  show  how  all  of  the  H2  and  N2  taken  may  be  converted 
into  NH3. 

524.  The  Effect  of  Removing  One  Product  of  a  Reaction.— 
In  the  discussion  of  the  reaction 

NaCl+H2SO4->NaHSO4+HCl  (289) 

it  was  shown  that  if  the  HC1  continuously  escapes  as  a  gas  the 
reaction  will  go  to  completion  from  left  to  right.  In  the  reaction 
between  hydrogen  and  nitrogen  equilibrium  results  from  two 
opposing  changes,  each  of  which  continues  in  the  equilibrium 
mixture :  these  are  the  union  of  H2  and  N2  on  the  one  hand  and 
the  dissociation  of  NH3  on  the  other.  If  the  NH3  formed  can 
be  removed  from  the  mixture  in  some  systematic  way,  the  H2 
and  N2  will  finally  be  completely  united.  We  shall  now  briefly 
describe  the  process  as  actually  carried  out. 

525.  The  Manufacture  of  Synthetic  Ammonia. — The  nitrogen 
used  in  making  synthetic  ammonia  might  be  obtained  from  the 
air  by  the  Linde  process  (513),  and  the  hydrogen  could  be  made 
by  the  electrolysis  of  water,  if  these  methods  were  not  too 
expensive .     A  practical  method  of  obtaining  the  required  mixture 
of  N2  and  H2  is  said  to  have  been  worked  out,  in  which  air  and 
steam  are  passed  over  heated  coke.     The  carbon  comprising  the 
coke  unites  with  the  oxygen  of  both  the  air  and  the  steam.     The 
CO2  and  CO  formed  are  then  removed,  leaving  N2  and  H2. 

The  apparatus  used  in  the  production  of  NH3  is  represented 
diagrammatically  in  Fig.  89.  The  mixture  of  N2  and  H2  under 
high  pressure,  say  100  atmospheres,  enters  at  A  and  passes  into 
the  chamber  B,  which  holds  the  catalytic  agent  and  is  surrounded 
by  an  electric  heater  C.  The  reaction 

3H2+N2->2NH3 

takes  place  to  the  extent  of  perhaps  8  to  10  per  cent  in  B,  and 
the  mixture  of  the  three  gases  passes  on  to  the  vessel  D,  sur- 
rounded by  the  refrigerating  vessel  E.  The  high  pressure  and 


332 


Introduction  to  General  Chemistry 


low  temperature  cause  nearly  all  of  the  NH3  to  condense  in  D 
to  the  liquid  state,  while  the  H2  and  N2,  being  much  less  readily 
liquefied,  pass  out  and  by  means  of  the  pump,  F,  are  conveyed 
by  a  pipe  back  to  the  reaction  vessel,  B.  Liquid  NH3  is  removed 
from  D  through  the  valve  at  the  bottom  of  D. 


FIG.  89 


526.  Ammonia    from    Cyanamid. — Cyanamid,    CaCN2,    is 
made  by  the  action  of  nitrogen  on  calcium  carbide  at  a  red  heat 

(514):  ' 

CaC2+N2->CaCN2+C. 

When  cyanamid  is  treated  with  steam  under  high  pressure 
the  following  reaction  takes  place: 

CaCN2+3H2O^2NH3+CaC03. 

Cyanamid  has  been  manufactured  at  Niagara  Falls,  Canada,  since 
1909;  and  since  1915  considerable  quantities  of  ammonia  have 
been  made  at  this  place  by  the  reaction  described  above. 
Recently  the  U.S.  government  has  built  an  immense  cyanide 
ammonia  plant  at  Muscle  Shoals,  Alabama.  It  now  appears 
possible  that  this  process  will  be  economically  superior  to 
that  by  which  ammonia  is  made  directly  from  nitrogen  and 
hydrogen. 

527.  The  Chemistry  of  Ammonia. — One  of  the  most  familiar 
reactions  of  ammonia  is  that  by  which  white  fumes  are  formed 


Nitrogen  and  Ammonia  333 

when  ammonia  gas  comes  in  contact  with  a  volatilized  acid. 
The  fumes  are,  of  course,  fine  particles  of  the  resulting  salt: 

NH3+HC1->NH4C1, 
NH3+HNO3-»NH4NO3. 

In  water  solution  we  have  the  same  type  of  reaction  with  acids, 
resulting  in  the  formation  of  ammonium  salts.  A  consideration 
of  the  equilibrium  equations 


will  show  that  we  may  represent  the  reaction  of  aqueous  ammonia 
solution  with  acids  as  the  result  of  the  union  of  the  hydrogen 
ion  of  the  acid  with  the  hydroxyl  ion  of  the  base  to  form  water, 
and  the  consequent  displacement  of  the  equilibrium  to  the  right; 
or  we  may  represent  the  reaction  as  the  union  of  NH3  with  H+ 
of  the  acid  to  form  NH4+,  and  the  consequent  displacement  of 
the  equilibrium  to  the  left.  The  net  result  of  either  reaction  is 
the  same,  and  very  likely  the  truth  is  that  both  routes  are 
actually  followed.  The  final  solution  contains  water  and  the 
highly  ionized  ammonium  salt. 

In  these  reactions  of  ammonia,  either  as  a  gas  or  in  solution, 
the  nitrogen  atom  suddenly  and  very  easily  takes  on  one  posi- 
tive and  one  negative  atom  or  radical.  So  striking  is  the  ease 
with  which  these  reactions  occur  that  we  think  that  the  nitrogen 
atom  was  ready  to  take  on  the  new  groups  and  so  must  carry 
a  positive  and  a  negative  charge,  balanced  against  each  other,  in 
addition  to  the  three  negative  charges  by  which  it  holds  the  three 
hydrogen  atoms,  thus,  =±=NH3.  From  what  we  know  of  the 
complicated  structure  of  an  atom  it  seems  plausible  to  assume 
the  existence  of  a  positive  and  a  negative  charge  on  the  same 
atom.  The  reaction  of  aqueous  ammonia  with  an  acid  may  then 
be  represented  as  follows: 


A  compound  which,  like  ammonia,  contains  an  atom  which  is 
not  exerting  its  full  valence  is  said  to  be  unsaturated.  There 
are  many  other  unsaturated  compounds. 


334  Introduction  to  General  Chemistry 

Though  the  above-mentioned  reactions  are  the  most  char- 
acteristic ones  shown  by  ammonia,  we  have  seen  one  case  of  a 
different  kind,  namely  the  reducing  action  of  ammonia  on  copper 

oxide  (52,  84): 

2NH3+3CuO-»N2+3H20-f-3Cu. 

An  even  more  vigorous  oxidation  of  ammonia  occurs  when  a 
mixture  of  ammonia  and  air  is  passed  over  red-hot  platinum. 
Under  the  catalytic  influence  of  platinum,  nitric  acid  is  formed: 

NH3+  2O2->HNO3+H2O. 

This  reaction  is  the  basis  of  one  of  the  great  processes  for  manu- 
facturing nitric  acid  from  the  air  (570). 

When  ammonia  gas  is  passed  over  heated  sodium  the  follow- 
ing reaction  takes  place: 

2Na+  2NH3-»2NaNH2+H2. 

The  products  are  hydrogen  and  sodium  amide.  The  latter  sub- 
stance is  completely  decomposed  by  water,  thus: 

NaNH2+H2O->NaOH+NH3. 

Ammonia  reacts  with  other  metals,  such  as  magnesium,  forming 
nitrides  (514),  compounds  which  are  also  hydrolyzed  by  water 
to  form  ammonia  and  the  hydroxide  of  the  metal. 

528.  Ammonium  Salts. — Ammonium  salts  closely  resemble 
the  corresponding  salts  of  potassium.    They  are  easily  soluble 
in  water  and  crystallize  well  from  solutions.    None  of  the  salts 
are  colored  excepting  those  with  colored  acid  ions. 

529.  Dissociation    of    Ammonium    Chloride. — Ammonium 
chloride  may  be  completely  volatilized  at  temperatures  much 
below  red  heat.    When  the  molecular  weight  of  the  salt  was 
calculated  on  the  basis  of  determinations  of  the  density  of  the 
vapor  formed  at  about  300°,  the  value  found  was  half  a  formula 
weight.    Apparently    a    full    formula    weight    of    ammonium 
chloride,  if  still  a  gas,  at  o°  and  76  cm.  pressure  would  occupy 
not  22.4  but  44.8  liters.    To  some  chemists  sixty  to  eighty 
years  ago  it  seemed  impossible  to  co-ordinate  these  facts  with 


Nitrogen  and  Ammonia  335 

Avogadro's  hypothesis,  and  they  accordingly  abandoned  the 
latter. 

A  simple  experiment  gives  the  clue,  however,  to  the  explana- 
tion of  the  enigma.  If  a  bit  of  dry  NH4C1  is  placed  in  a  dry  test 
tube  and  cautiously  heated,  a  piece  of  moist  red  litmus  paper 
held  in  the  mouth  of  the  tube  turns  blue,  thus  indicating  the 
formation  of  free  NH3.  Elaborate  experiments  have  shown 
that  the  vapor  of  NH4C1  dissociates  at  300°  almost  completely 


into  NH3  and  HC1: 


NH4C1±*NH3+HCL 


Since  the  molecules  of  ammonia  are  much  lighter  than  those  of 
hydrogen  chloride  they  travel  (diffuse)  more  rapidly  at  the  same 
temperature  (197).  Hence,  although  the  two  gases  are  released 
at  the  same  instant,  the  ammonia  molecules  are  the  first  to  reach 
the  litmus  paper  in  the  mouth  of  the  test  tube.  Upon  cooling, 
the  constituents  unite  again  to  form  solid  NH4C1.  Therefore 
NH4C1  vapor  has  half  the  density  expected,  for  the  reason  that 
the  number  of  molecules  present  is  practically  double  the 
original,  because  each  NH4C1  molecule  gives  one  molecule  of 
NH3  and  one  of  HC1.  Thus  Avogadro's  law  covers  the  case 
perfectly.  Many  other  apparent  exceptions  to  Avogadro's  law 
are  known,  but  in  every  case  a  satisfactory  explanation  has  been 
found  in  complete  harmony  with  the  law. 

530.  Dissociation  of  Other  Ammonium  Salts. — Several  other 
ammonium  salts  also  dissociate  into  the  acid  and  ammonia. 
Thus  solid  ammonium  bicarbonate  smells  strongly  of  ammonia, 
at  ordinary  temperatures,  by  reason  of  the  following  dissociation: 

NH4HCO3^NH3-fH2O+CO2. 

Ammonium  sulfide,  formed  by  the  union  of  ammonia  with 
hydrogen  sulfide,  dissociates  at  room  temperatures  to  form 
ammonia  and  hydrogen  sulfide: 

(NH4)HS^NH3+H2S. 

Ammonium  hydroxide  itself,  which  may  be  considered  as  the 
product  of  the  union  of  ammonia  and  the  very  weak  acid  water, 


336  Introduction  to  General  Chemistry 

is  stable  in  crystalline  form  only  at  a  very  low  temperature  (91). 
In  general  we  find  that  the  weaker  the  acid  the  more  unstable 
is  the  ammonium  salt.  It  is  perhaps  well  to  point  out  that  dis- 
sociations of  the  kind  just  considered  are  entirely  distinct  from 
the  splitting  up  of  salts,  acids,  and  bases  into  ions  in  solution,  in 
that  the  new  molecules  formed  by  the  dissociation  of  a  gas  or 
vapor  are  not  electrically  charged.  Furthermore  the  dissociation 
products  are  of  different  composition  in  the  two  cases,  as  is 
illustrated  by  Table  XXI. 

TABLE  XXI 


Substance 

Electrolytic 
Dissociation 

Gaseous 
Dissociation 

NH4C1.. 
NH4HCO3.... 
NH4HS 

NH.+      Cl- 
NH4+      HCO3- 
NH4+      HS- 

NH3      HC1 
NH3      H2O      CO2 
NH3      H2S 

NH4OH  

NH4+      OH- 

NH3      H20 

531.  Other  Compounds  of  Nitrogen  and  Hydrogen. — Nitro- 
gen and  hydrogen  form  several  other  compounds  besides  NH3; 
among  such  hydrazine,  N2H4,  and  its  hydrate,  N2HSOH,  a  basic 
substance,  may  be  mentioned.    Hydroxylamine,   NH2OH,   is 
also  a  base.    Hydrazoic  acid,  HN3,  is  a  violently  explosive 
liquid.     It  is  an  acid  and  forms  salts  like  NaN3,  sodium  azide, 
and  AgN3,  silver  azide;  the  latter  is  extremely  explosive.    The 
ammonium  salt  NH4N3  and  the  hydrazine  salt  N2H5N3  may  be 
considered  as  having  the  formulae  N4H4  and  N5HS  respectively. 

532.  The  Solubility  of  Silver  Chloride  in  Ammonia  Solu- 
tion.— Silver  chloride,  AgCl,  as  we  have  seen  (169),  is  an  almost 
insoluble  white  salt  which  is  easily  obtained  by  adding  any 
chloride  solution  to  a  silver  salt  solution;   for  example, 

NaCl+AgNO3->AgCl-hNaNO3. 

This  reaction  goes  nearly  to  completion  by  reason  of  the  very 
small  solubility  of  silver  chloride,  so  that  precipitation  con- 
tinues until  the  concentrations  of  both  Ag+  and  Cl~  ions  are 
very  small  (452). 


Nitrogen  and  Ammonia  337 

If  pure  silver  chloride  is  stirred  with  water,  only  very  little 
dissolves;  in  fact,  about  i  .  3  mg.  per  liter.    We  have,  then, 

Solid  Dissolved 


By  far  the  greater  part  of  the  dissolved  substance  is  present  as 
Ag+  and  Cl~  ions. 

If  a  solution  of  ammonia  is  added  to  the  AgCl  and  water  the 
solid  dissolves  very  easily,  giving  a  clear,  colorless  solution. 
This  solution  is  a  good  electrical  conductor,  while,  as  we  have 
already  learned,  ammonium  hydroxide  conducts  very  poorly. 
These  facts  may  be  demonstrated  by  the  use  of  the  apparatus 
described  in  sec.  384.  If  the  cell  is  first  filled  with  ammonia 
solution,  it  is  found  that  the  current  which  passes  is  very  weak. 
As  soon  as  AgCl  is  dissolved  in  the  ammonia  solution  the  galva- 
nometer needle  is  strongly  deflected.  As  a  matter  of  fact  the 
new  solution  is  as  good  a  conductor  as  solutions  of  equal  con- 
centrations of  most  salts. 

533.  The  Silver  Ammonium  Ion,  Ag(NH3)2+.  —  It  has  long 
been  known  that  dry  AgCl  unites  with  dry  NH3  gas  to  form 
solid  compounds.  It  is  possible  that  one  or  more  of  these  is 
present  in  the  solution  formed  when  AgCl  dissolves  in  a  solution 
of  ammonia;  but  if  so,  what  are  the  ions,  if  any,  of  the  new  silver 
compounds?  We  have  already  learned  (397)  that  the  composi- 
tion of  the  ions  of  a  solution  is  best  investigated  by  means  of 
migration  experiments.  If  we  electrolyze  a  solution  of  AgCl 
dissolved  in  ammonia,  using  a  U-tube  like  that  shown  in  Fig.  47 
(397)>  we  find  that  both  Ag+  and  NH3  migrate  toward  the 
cathode  or  negative  electrode,  while  the  chloride  ion  migrates  in 
the  opposite  direction.  Moreover,  we  find  that  the  Ag+  and 
NH3  migrating  toward  the  cathode  are  in  the  proportion  of  one 
atom  of  Ag  to  two  molecules  of  NH3  as  represented  in  the 
formula  Ag(NH3)2,  so  that  it  would  appear  that  this  last  formula 
represents  the  composition  of  the  positive  ion.  We  are  therefore 
led  to  conclude  that  the  following  reaction  first  takes  place, 

AgCl+2NH3±,Ag(NH3)2Cl, 


Introduction  to  General  Chemistry 

and  that  the  compound  so  formed  ionizes  very  readily  thus  : 

Ag(NH3)2Cl^Ag(NH3)2++Cl-u 

The  silver  ammonium  ion,'Ag(NH3)2+,  is  called  a  complex  ion. 

534.  The  Stability  of  the  Ag(NH3)2+  Ion.—  It  has  been  proved, 
by  an  electrical  method  which  need  not  be  considered  here,  that 
in  a  solution  of  silver  chloride  in  ammonia  only  an  exceedingly 
small  number  -  of  Ag+  ions  are  present.    This  means  that  the 
complex  ion,  Ag(NH3)2+,  is  very  stable  and  is  dissociated  only 
very  slightly  into  Ag+  and  2NH3,  thus: 

Ag(NH3)2+^Ag++2NH3. 

535.  Why  Silver  Chloride  Dissolves  in  Ammonia.  —  We  can 

now  explain  why  AgCl  dissolves  in  ammonia  solution.     Let  us 
consider  the  following  formulation  of  the  reaction: 

Solid  Dissolved 


2NH3 

It 
Ag(NH3)2+. 

The  larger  part  of  the  silver  chloride  dissolved  by  pure  water  is 
present  as  Ag+  and  Cl~  ions.  If  NH3  is  added,  the  Ag+  ions 
unite  with  it  nearly  completely  to  form  Ag(NH3)2+.  This 
reaction  greatly  reduces  the  concentration  of  Ag+  and  therefore 
causes  more  solid  AgCl  to  pass  into  solution.  If  sufficient  NH3 
is  present  these  changes  go  on  until  all  the  solid  AgCl  has  dis- 
solved. In  the  solution  we  have  the  salt  silver  ammonium 
chloride. 

Consideration  of  the  equation 

Ag++2NH3^Ag(NH3)2+ 

shows  that  by  adding  an  excess  of  NH3  the  concentration  of  Ag  *" 
will  be  still  further  decreased.  If  we  wish  to  dissolve  AgCl  com- 
pletely it  is  necessary  to  add  a  little  more  than  two  formula 
weights  of  NH3  for  one  of  AgCl  in  order  to  depress  the  con- 
centration of  Ag+  sufficiently.  The  dissolving  of  silver  chloride 


Nitrogen  and  Ammonia  339 

in  ammonia  is  exactly  similar  to  the  dissolving  of  silver  acetate 
in  nitric  acid,  except  that  in  the  latter  case  the  acetate  ion  was 
suppressed,  while  in  the  former  the  Ag+  ion  was  suppressed. 

536.  The  Effect  of  Acids  on  Silver  Ammonium  Ion.—  If  a 
strong  acid,  nitric  acid  for  example,  is  added  to  silver  ammo- 
nium chloride  solution,  silver  chloride  appears  as  a  precipitate. 
Evidently  the  nitric  acid  has  caused  an  increase  in  the  con- 
centration of  Ag+  ion,  and  union  of  the  latter  with  the  abund- 
ant chloride  ion  has  followed.  This  reaction  is  explained  by 
saying  that  NH3  is  converted  into  NH4NO3  by  the  HN03. 
This  the  acid  does  by  virtue  of  the  hydrogen  ion  which  it 
furnishes.  The  latter  unites  with  the  free  ammonia  as  follows: 


Some  union  of  the  latter  ion  with  the  nitrate  ion  to  form  NH4N03 
also  takes  place,  but  the  critical  change  is  that  of  ammonia  to 
ammonium  ion.  The  removal  of  NH3  allows  the  dissociation  of 
the  silver  ammonium  ion  to  proceed  to  completion: 

Ag(NH3)2+->Ag++2NH3. 

537.  Other  Silver  Ammonium  Compounds.  —  If  ammonium 
hydroxide  is  added  to  silver  nitrate  solution  a  little  at  a  time, 
brown  silver  hydroxide  is  seen  to  precipitate  but  eventually  to 
redissolve,  just  as  the  ammonia  added  exceeds  the  proportion 
of  two  molecules  for  each  molecule  of  silver  nitrate  present  : 

AgNO3+NH4OH->NH4NO3+AgOH; 
AgOH+  2NH3->Ag(NH3)2OH. 

But  silver  ammonium  hydroxide  is  a  strong  base,  and  as  fast  as  it 
forms  it  displaces  ammonium  hydroxide  from  its  salt,  NH4N03J 

NH4NO3+Ag(NH3)3OH->NH4OH-fAg(NH3)2N03, 

so  that  the  final  equation  for  the  formation  of  silver  ammonium 
nitrate  from  ammonia  and  silver  nitrate  is 

AgN03+  2NH3-»Ag(NH3)2N03. 


340  Introduction  to  General  Chemistry 

In  the  same  way  a  solution  of  silver  ammonium  sulfate  can 
be  made  by  the  addition  of  aqueous  ammonia  to  silver  sulfate. 
All  the  silver  ammonium  salts  are  found  to  be  highly  ionized, 
just  as  are  potassium  and  ammonium  salts. 

538.  Other  Complex  Ions. — Ammonia  unites  with  many 
other  metal  ions  to  form  complex  ions;  for  instance,  with  cupric 
ion  it  forms  the  brilliant  blue  copper  ammonium  ion,  and 
with  nickel  ion  it  forms  the  less  highly  colored  blue  nickel 

ammonium  ion. 

Cu+++4NH3->Cu(NH3)4++ 

Ni++ +4NH3-^Ni(NH3)4++ . 

These  complex  ions  form  salts  such  as  Cu(NH3)4S04  and 
Ni(NH3)4(NO3)2.  The  chemistry  of  these  ions  parallels  that  of 
the  silver  ammonium  ion.  The  ammonium  ions  are  the  first 
which  we  have  called  complex  ions.  But  as  a  matter  of  fact 
we  have  been  dealing  with  a  great  number,  for  example,  nitrate, 
carbonate,  phosphate,  sulfate,  and  hydroxyl  ions.  In  the  double 
decomposition  reactions  studied,  these  ions  have  shown  no  sign 
of  any  dissociation  into  smaller  parts  corresponding  to  the 
dissociation  of  the  silver  ammonium  ion  into  silver  ion  and 
ammonia.  However,  there  is  evidence  that  these  secondary 
dissociations  do  exist,  but  to  a  very  much  smaller  degree  than 
in  the  case  of  a  complex  silver  ion.  In  other  words,  these  ions 
are  very  much  more  stable  than  the  latter.  We  shall  find  later 
other  examples  of  moderately  stable-  complex  ions  like  those  of 
ammonia. 


CHAPTER  XXII 
NITRIC  ACID  AND  THE  OXIDES  OF  NITROGEN 

539.  Nitric  Acid,  HNO3. — That  nitric  acid  is  a  substance  of 
great  importance  will  be  apparent  at  once  when  it  is  known  that 
it  is  an  indispensable  agent  in  the  manufacture  of  explosives  and 
dyestuffs  and  in  addition  is  used  extensively  in  a  great  many 
other  ways.     We  have  already  learned  (104)  that  nitric  acid  can 
be  made  from  sodium  nitrate,  Chile  saltpeter,  and  sulfuric  acid : 

NaNO3+H2SO4->NaHSO4+HNO3. 

If  a  larger  proportion  of  NaNO3  is  used,  the  following  reaction 
can  also  occur  if  the  temperature  is  high  enough: 

2NaNO3-f  H2SO4->Na2SO4+  2HNO3. 

Until  recently  the  only  commercial  method  of  making  nitric 
acid  was  by  means  of  these  two  reactions.  In  practice  the 
action  takes  place  in  large  cast-iron  stills,  which  are  but  slightly 
attacked  by  the  two  acids  as  long  as  water  is  not  present.  The 
nitric  acid  distils  and  is  usually  condensed  in  vessels  made  of 
fused  quartz,  which  is  not  acted  upon  by  HNO3,  or  in  Duriron  (an 
alloy  of  iron  and  silicon),  which  is  almost  unaffected  by  the  acid. 

540.  Chile  Saltpeter,  NaNO3.— It  is  a  remarkable  fact  that 
there  is  but  one  known  source  of  sodium  nitrate  of  sufficient 
magnitude  to  be  of  practical  importance:  this  consists  of  enor- 
mous deposits  found  in  a  desert  region  in  the  mountains  of  Chile. 
These  deposits  form  a  stratum  averaging  five  feet  in  thickness 
and  covering  about  six  hundred  square  miles.    The  value  of  the 
saltpeter  exported  amounts  to  three-fourths  of  the  total  exports 
of  Chile,  and  the  export  duty  on  it  is  the  chief  source  of  revenue 
of  the  country. 

Potassium  nitrate  is  found  in  limited  quantities  in  India,  but 
the  deposits  are  far  too  small  to  be  of  importance  for  the  manu- 
facture of  nitric  acid. 

341 


342  Introduction  to  General  Chemistry 

541.  Properties  of   Nitric  Acid. — Pure  anhydrous   (water- 
free)  nitric  acid  is  a  colorless  liquid  of  density  1.52;  it  boils  at 
78°,  undergoing  some  decomposition  into  water  and  oxides  of 
nitrogen.    The  pure  acid  of  commerce  contains  only  68  per  cent 
of  HN03,  the  balance  being  water.    This  acid  has  a  density  of 
1.42.    When  dilute  nitric  acid  is  boiled  the  residual  portion 
grows  more  concentrated,  until  it  reaches  a  density  of  i .  42  and 
68  per  cent  purity.    Acid  of  this  concentration  then  distils  at 
the  constant  temperature  of   120°.     Acid  more  concentrated 
than  68  per  cent  boils  at  lower  temperatures,  and  as  the  boil- 
ing continues  the   temperature  rises,  while  the  residual  acid 
becomes  less  concentrated,  until  finally,  at  120°,  68  per  cent 
acid  is  left. 

Nitric  acid  unites  with  water  to  form  two  crystalline  hydrates, 
HN03-H2O  and  HN03-3H20,  which  melt  at  -38°  and  -18° 
respectively. 

Nitric  acid  is  one  of  the  most  highly  ionized  acids;  it  therefore 
has  all  the  properties  characteristic  of  a  strong  acid.  Its  salts 
are  all  soluble  in  water — most  of  them  very  soluble.  Nitrates 
of  strong  bases  (Na,  K,  Ca,  Ba,  Mg,  and  Ag)  give  neutral  solu- 
tions, while  solutions  of  nitrates  of  weak  bases  are  acid  in  reaction, 
owing  to  hydrolysis  (436).  Other  important  properties  demand 
treatment  in  separate  paragraphs. 

542.  Nitric  Acid  as  an  Oxidizing  Agent. — In  addition  to  its 
action  as  an  acid,  nitric  acid  also  acts  as  a  powerful  oxidizing 
agent.    It  will  be  recalled  that  oxidation  and  reduction  always 
go  hand  in  hand  (327) ;  the  oxidizing  agent  is  reduced,  and  the 
reducing  agent  is  oxidized.     Most  reducing  agents  react  readily 
with  nitric  acid,  being  themselves  oxidized  as  the  result  of  the 
reaction.    The  reduction  of  HNO3  gives  the  compounds  N02, 
NO,  N20,  N2,  or  NH3,  according  to  the  reducing  agent  used. 
The  subject  is  best  approached  after  one  has  become  familiar 
with  the  properties  of  the  oxides.    We  shall  therefore  first  take 
up  the  study  of  the  latter  and  then  return  to  the  discussion  of 
the  action  of  nitric  acid  as  an  oxidizing  agent  (557). 

543.  Nitric   Oxide,   NO. — Nitric   oxide  is  a  colorless  gas 
almost  insoluble  in  water;    it  is  most  easily  made  by  the 


Nitric  Acid  and  Oxides  of  Nitrogen  343 

action  of  dilute  HN03  on   copper.     We  might  expect   the 

reaction  to  be 

Cu+  2HN03->Cu(N03)2+H2, 

but  copper  has  virtually  no  tendency  to  displace  hydrogen  from 
solutions  of  acids.  We  find  that  the  reaction  actually  takes 
place  according  to  the  equation 

3Cu+8HN03->3Cu(N03)2+4H20+2NO. 

The  apparatus  shown  in  Fig.  87  (513)  is  used.  About  20  g.  of 
copper  in  the  form  of  wire,  clippings,  turnings,  etc.,  are  placed 
in  a  250  c.c.  flask,  and  60  c.c.  of  concentrated  nitric  acid  mixed 
with  60  c.c.  of  water  are  added.  The  wash  bottle  contains  dilute 
NaOH. 

The  reaction  in  the  flask  starts  at  once,  and  the  vessel  is  soon 
filled  with  a  brown-colored  gas.  After  a  few  minutes  this  dis- 
appears, and  in  its  place  is  a  faintly  colored  gas,  nitric  oxide, 
which  is  impure.  When  purified  by  passing  through  the  wash 
bottle  it  proves  to  be  colorless.  It  may  now  be  collected  over 
water  in  the  cylinder  set  in  the  pneumatic  trough. 

544.  Direct  Union  of  Nitrogen  and  Oxygen. — The  student's 
general  chemical  experience  and  observations  will  have  already 
led  him  to  the  conclusion  that  nitrogen  does  not  readily  combine 
directly  with  oxygen;  for  if  the  reaction 

N2+O2->2NO 

occurred  very  easily  it  would  be  impossible  for  the  mixture  of 
N2  and  O2  composing  air  to  exist  without  union  taking  place. 
Nevertheless  the  passage  of  electric  sparks  through  air  leads  to 
the  formation  of  a  small  amount  of  nitric  oxide;  but  the  reaction 
soon  reaches  equilibrium  because  of  the  decomposition  of  nitric 
oxide  into  its  constituents  under  the  action  of  electric  sparks. 
The  passage  of  a  flash  of  lightning  through  air  produces  a  little 
nitric  oxide. 

Although  the  union  of  nitrogen  and  oxygen  under  the  influ- 
ence of  electric  sparks  was  discovered  by  Cavendish  in  1766,  it 
had  only  a  scientific  interest  for  chemists  up  to  very  recent  times. 


344  Introduction  to  General  Chemistry 

The  modern  application  of  this  reaction  to  the  manufacture  of 
nitric  acid  will  be  taken  up  later  (566). 

545.  The  Properties  of  Nitric  Oxide,  NO. — Nitric  oxide  is  a 
colorless  gas  nearly  insoluble  in  water  and  slightly  heavier  than 
air.    One  liter  weighs  i .  34  g.  and  22.4  liters  30  g.,  agreeing  with 
that  calculated  from  the  formula  NO.    Nitric  oxide  is  very 
stable,  not  being  decomposed  into  its  elements  except  at  high 
temperatures. 

When  nitric  oxide  comes  in  contact  with  air  it  turns  brown. 
The  same  change  takes  place  when  nitric  oxide  is  mixed  with  half 
its  volume  of  pure  oxygen.  The  reaction  takes  place  thus : 

2NO+O2->2NO2. 

The  brown  gas,  N02,  is  called  nitrogen  tetroxide.  This  is  the 
brown  gas  that  first  appears  when  nitric  oxide  is  being  made, 
since  the  latter  unites  with  the  O2  present  in  the  flask  to  form 
brown  N02.  It  is  interesting  to  note  that  a  little  heat  is  given 
out  when  NO  and  O2  unite. 

546.  Combustion  in  Nitric  Oxide. — The  proportion  of  oxygen 
in  nitric  oxide  is  about  two  and  one-half  times  as  great  as  in  air? 
and  we  might  expect  that  this  gas  would  support  combustion  very 
readily;  but  this  is  not  the  case.     Burning  sulfur  or  a  burning 
candle  is  extinguished  if  brought  into  a  jar  of  nitric  oxide.    A 
mixture  of  hydrogen  and  nitric  oxide  does  not  explode  with  an 
electric  spark.    On  the  other  hand,  a  few  substances  will  burn 
in  nitric  oxide  if  strongly  heated;   thus  an  iron  wire  heated  to 
incandescence  by  an  electric  current  takes  fire  in  this  gas  and 
burns,  forming  an  oxide  and  setting  nitrogen  free.    Briskly 
burning  phosphorus  continues  to  burn  when  plunged  into  a  jar 
of  nitric  oxide,  forming  phosphorous  pentoxide  and  nitrogen, 

4?+ ioNO-»2P2Os+ sN2. 

The  vapor  of  carbon  bisulfide,  CS2,  an  easily  combustible 
and  very  volatile  liquid,  gives  with  nitric  oxide  a  mixture  that 
readily  burns  with  a  bright  light.  By  way  of  caution  it  may  be 
pointed  out  that  carbon  bisulfide  is  a  very  dangerous  liquid, 


Nitric  Acid  and  Oxides  of  Nitrogen  345 

because  of  the  readiness  with  which  it  catches  fire.    It  must  be  kept 
away  from  all  flames  and  heated  objects. 

547.  Other  Ways  of  Making  Nitric  Oxide.  —  Nitric  oxide  is 
formed  by  the  action  of  dilute  nitric  acid  on  many  other  metals 
besides  copper.     It  is  also  formed  when  various  oxidizable  sub- 
stances react  with  dilute  HNO3  as  illustrated  by  the  following 
equation  : 

2HNO3+6FeSO4-f3H2S04->3Fe2(SO4)3-f2NO+4H20. 

By  means  of  this  reaction  very  pure  nitric  oxide  can  be  made. 

548.  Nitric  Oxide  and  Ferrous  Sulfate.  —  When  nitric  oxide 
is  passed  into  a  solution  of  FeS04  a  dark-colored  solution  is 
formed  containing  the  compound  FeS04-NO.    When  this  solu- 
tion is  heated  the  compound  is  decomposed,  and  very  pure 
nitric  oxide  is  given  off.     The  formation  of  a  nearly  black  solu- 
tion with  FeSO4  serves  as  a  good  test  for  nitric  oxide,  and  in  con- 
sequence also  for  nitric  acid,  by  reason  of  the  fact  that  nitric 
acid  is  reduced  to  nitric  oxide  by  FeS04  in  the  presence  of  H2S04, 
the  ferrous  sulfate  being  oxidized  to  ferric  sulfate,  Fe2(S04)3,  at 
the  same  time. 

549.  A  Test  for  Nitric  Acid  and  Nitrates.  —  If  2  or  3  c.c.  of  a 
mixture  of  nitric  acid  or  any  nitrate  and  concentrated  sulfuric 
acid  are  placed  in  the  bottom  of  a  test  tube,  and  a  solution  of 
ferrous  sulfate  is  cautiously  added  in  such  a  way  that  the  two 
solutions  do  not  mix,  a  dark-brown  ring  will  be  formed  at  the 
junction  of  the  two  layers.    The  brown  layer  contains  the  com- 
pound FeS04-  NO. 

550.  Nitrogen  Tetroxide,  NO2.—  We  have  learned  that  the 
brown  gas  nitrogen  tetroxide  is  formed  by  the  direct  union  of 
nitric  oxide  and  oxygen,  thus: 


It  is  also  formed  in  several  other  ways,  such  as  by  the  action  of 
concentrated  HN03  on  many  metals.  For  example,  with  con- 
centrated HNO3  and  copper  we  have 

Cu+4HNO3->Cu(NO3)2-{-2NO2-f2HaO. 


346  Introduction  to  General  Chemistry 

Dilute  HNO3  and  copper  give  NO  instead  of  N02  (543). 
The  cause  of  this  difference  is  found  in  the  behavior  of  the  two 
oxides  of  nitrogen  toward  H2O  and  HNO3,  as  represented  in  the 
equation 


With  much  warm  water  this  reaction  takes  place  nearly  com- 
pletely from  left  to  right;  but  the  reaction  is  reversible,  since 
nitric  oxide  acts  almost  completely  on  an  excess  of  concentrated 
nitric  acid  to  form  nitrogen  tetroxide  and  water.  Therefore,  in 
the  action  of  nitric  acid  on  copper,  nitric  oxide  will  be  given  off 
if  the  acid  is  very  dilute,  and  nitrogen  tetroxide  if  it  is  very  con- 
centrated; with  an  acid  of  intermediate  concentration  a  mixture 
of  the  two  oxides  will  result. 

Nitrogen  tetroxide  is  also  prepared  by  heating  the  nitrates 
of  heavy  metals  such  as  copper,  lead,  mercury,  silver,  etc.  (565). 

551.  The   Physical    Properties    of    Nitrogen    Tetroxide.— 
Nitrogen  tetroxide  is  a  reddish-brown  gas  having  a  peculiar, 
disagreeable  odor.     It  is  dangerously  poisonous  and  may  pro- 
duce fatal  results  after  a  day  or  two  when  inhaled  in  quantities 
which  seem  inconsequential  at  the  time.    It  is  easily  condensed 
to  liquid  form  by  cooling  it  with  ice;  at  —  10°  liquid  nitrogen 
tetroxide  solidifies  to  a  mass  of  nearly  colorless  crystals.    The 
liquid  formed  by  the  melting  of  these  crystals  is  also  nearly 
colorless  at  temperatures  near  the  melting-point  ;  as  the  tempera- 
ture is  raised  the  liquid  becomes  first  yellow,  then  orange,  and 
finally  boils  at  22°,  giving  a  gas  or  vapor  of  light  reddish-brown 
color.    At  higher  temperatures  the  gas  becomes  much  darker  in 
color. 

552.  The  Density  of  Nitrogen  Tetroxide.  —  The  formula  N02 
corresponds  to  a  formula  or  molecular  weight  of  46  and  a  gas 
density  i  .  59  times  that  of  air.     The  actual  density  which  at 
room  temperature  is  very  much  greater  than  this  decreases 
with  rise  of  temperature  and  reaches  the  value  i  .  59  only  at 
140°.    The  exact  results  are  shown  in    Table  XXII. 

The  very  remarkable  changes  in  density  and  color  with 
change  of  temperature  are  thought  to  be  due  to  the  formation 


Nitric  Acid  and  Oxides  of  Nitrogen 


347 


at  low  temperatures  of  double  molecules  having  the  formula 
N204,  the  equation  being 


If  we  accept  this  view  and  also  suppose  N2O4  to  be  colorless  and 
NO2  dark  reddish-brown  we  get  a  very  satisfactory  explanation  of 
all  the  known  facts.  The  form  N204  exists  almost  pure  in  the 
crystals  at  —  10°;  in  the  liquid  state  some  dissociation  to  NO2 
has  taken  place,  giving  a  yellow  or  at  higher  temperature  an 
orange-colored  liquid. 

TABLE  XXII 


Density 

Calculated 

Temperature 

Air  =  i 

Molecular  Weight 

Per  Cent  of  NOi 

Per  Cent  of  NdO« 

25°  

2.50 

72                              43 

57 

97°  

1.78                            52                              87 

13 

140°  

1-59 

46 

IOO 

0 

In  Table  XXII  the  percentages  of  NO2  and  N204  at  various 
temperatures  are  given,  as  calculated  from  the  densities.  At 
140°  the  dissociation  into  single  molecules  of  NO2  is  practically 
complete.  In  accord  with  this  explanation  it  is  actually  found 
that  the  density  (corrected  of  course  for  thermal  expansion) 
remains  constant  above  140°. 

In  chemical  equations  we  shall  continue  to  write  the  formula 
N02  for  nitrogen  tetroxide,  but  it  will  be  understood  that  much 
of  the  gas,  at  room  temperature,  is  in  the  form  of  N204. 

553.  Nitrites  and  Nitrous  Acid. — A  very  interesting  reaction 
takes  place  when  N02  is  passed  into  a  solution  of  NaOH.  The 
gas  is  completely  absorbed,  and  the  solution  upon  evaporation 
yields  crystals  of  two  salts — sodium  nitrate,  NaN03,  and  sodium 
nitrite,  NaN02.  The  reaction  occurs  according  to  the  equation 

2NO2+  2NaOH-*NaNO3-f  NaNO2-f  H2O. 

Sodium  nitrite  is  a  salt  of  technical  importance,  particularly 
in  connection  with  the  manufacture  of  certain  dyestuffs  and 
medicinal  chemicals.  It  is  the  salt  of  an  acid,  HNO2,  called 


34&  Introduction  to  General  Chemistry 

nitrous  acid.  The  free  acid  itself  is  very  unstable  and  its  dilute 
solution  easily  decomposes,  thus: 

3HN02->HN03+  2NO+H20. 

On  the  other  hand  the  nitrites  of  sodium  and  potassium  are 
stable  in  water  solution.  They  are  also  resistant  to  heat.  As 
a  matter  of  fact  they  are  usually  made  by  the  decomposition 
of  the  corresponding  nitrates  under  heat  (565).  Ammonium 
nitrite,  however,  decomposes  in  hot-water  solution,  forming 
nitrogen  and  water  (513). 

All  nitrites  are  easily  oxidized  to  nitrates.  Solutions  of  these 
reagents  undergo  this  change  simply  on  exposure  to  air  and 
consequently  are  never  free  from  nitrates. 

Nitrous  acid  is  an  oxidizer  strong  enough  to  change  hydriodic 
acid  into  iodine  and  water,  but  not  strong  enough  to  act  similarly 
with  hydrobromic  acid: 

2HI+  2HNO2->I2+  2NO+  2H2O. 

Since  nitrous  acid  is  easily  reduced  to  nitric  oxide  it  will  give 
the  ring  test  (549),  as  the  latter  depends  on  the  formation  of 
this  substance. 

554.  Nitrogen  Trioxide,  or  Nitrous  Acid  Anhydride. — An 
oxide  having  the  composition  N203  is  found  to  form  white  crystals 
at  — 103°.    At  room  temperatures  it  can  exist  only  in  minimal 
amounts  in  equilibrium  with  its  decomposition  products  NO 

andN02: 

N203->NO+N02. 

The  decomposition  of  nitrous  acid  yields  products  which  appar- 
ently come  from  the  dissociation  of  this  gas,  so  that  the  latter 
may  be  called  the  anhydride  of  nitrous  acid  (313) : 

2HNO2^H2O-fN2O3. 

555.  Nitrogen  Pentoxide,  the  Anhydride  of  Nitric  Acid.— 
Nitrogen  pentoxide,  N2OS,  is  a  rather  unstable,  colorless  crystal- 
line substance  which  is  formed  from  pure  anhydrous  (water-free) 
HN03  by  the  action  of  P2OS,  thus: 

2HNO3+P2OS->2HPO3+N2OS. 


Nitric  Acid  and  Oxides  of  Nitrogen  349 

The  N2OS  formed  is  separated  from  the  involatile  metaphosphoric 
acid,  HPO3,  by  careful  distillation  and  condensation  in  a  vessel 
cooled  with  ice.  Nitrogen  pentoxide  decomposes  readily  into 
NO2  and  O2,  thus: 


It  also  unites  vigorously  with  H2O  to  form  HNO3, 


556.  Nitrous  Oxide,  N2O,  Laughing  Gas.  —  Of  the  five  oxides 
of  nitrogen,  nitrous  oxide,  N2O,  is  the  only  one  met  with  outside 
of  chemical  works  and  laboratories.     It  is  the  well-known  sub- 
stance laughing  gas,  used  by  all  dentists  as  a  mild  anaesthetic. 
This  gas  is   made  practically  by  heating  ammonium  nitrate, 
which  decomposes  easily  according  to  the  following  equation: 

NH4N03->N20+2H2O. 

If  the  substance  is  heated  too  strongly  the  reaction  takes  place 
explosively.  In  fact,  NH4NO3  is  being  used  extensively  as  an 
explosive.  Therefore  great  care  must  be  taken  to  heat  the  substance 
cautiously. 

Nitrous  oxide  is  a  colorless  gas,  having  a  faint,  not  unpleasant 
odor.  It  is  somewhat  soluble  in  water  but  is  usually  collected 
over  warm  water.  The  gas  is  condensed  to  liquid  form  by 
sufficient  pressure.  Liquid  N2O  boils  at  —90°.  This  liquid, 
contained  in  steel  cylinders,  is  an  article  of  commerce. 

Although  N20  contains  a  smaller  proportion  of  oxygen  than 
any  other  oxide  of  nitrogen,  it  supports  combustion  almost  as 
well  as  oxygen  itself.  A  glowing  splint  bursts  into  flame  when 
brought  into  N2O,  just  as  it  would  do  in  O2. 

557.  Oxidation  by  Nitric  Acid.  —  We  are  now  ready  to  discuss 
oxidation  by  nitric  acid.     Since  these  reactions  involve  changes  in 
electric  charges  on  the  nitrogen  atom  (501),  let  us  go  over  the 
common  reduction  products  of  nitric  acid  with  a  view  to  assign- 
ing the  electric  charges  to  the  nitrogen  atoms  present  in  them. 
In  doing  this  two  rules  are  followed:  (i)  oxygen  in  combination 
is  double  negatively  charged  except  in  the  case  of  a  peroxide; 


350  Introduction  to  General  Chemistry 

and  (2)  hydrogen  in  combination  is  always  single  positively 
charged.  Since  the  sum  of  the  charges  carried  by  the  atoms  in 
an  electrically  neutral  substance  must  be  zero,  we  may  calculate 
the  charge  on  one  of  the  atoms  of  a  compound  if  its  charge  is  the 
only  one  not  known.  Table  XXIII  shows  the  probable  charges 
on  the  nitrogen  atom  in  the  compounds  we  have  studied.  For 
convenience,  charges  greater  than  three  in  number  are  written 
with  a  superscript  to  show  the  number  of  charges.  Thus  N5"1" 
is  equivalent  to  N+++++.  When  nitric  acid  is  reduced  to 
NO2  the  original  atom  of  nitrogen  has  gained  one  electron 
(NS+^N4+).  To  form  N2  it  must  gain  five  electrons,  and  to 
form  NH3  it  must  gain  eight.  Therefore  it  is  a  greater  feat  to 


TABLE  XXIII 

Compounds 

HN03  NA  and  nitrates 

Nitrogen  Atoms 

NO2  NA 

N4+ 

HN02,  N2O3  and  nitrites    . 
NO     

.     N+++ 
.     .-   N++ 

N2O    

.     .     N+ 

N2  

.     .     N° 

NIL 

N  

reduce  nitric  acid  to  ammonia  than  to  nitric  oxide.  And  in 
general  we  find  that  the  more  active  the  reducing  agent  the  more 
we  get  of  products  like  N2O,  N2,  and  NH3.  Thus  the  main 
products  of  the  reaction  of  copper  and  nitric  acid  are  NO  and 
N02  (550)  ;  but  when  zinc  reacts  with  nitric  acid,  N2O,  N2,  and 
NH3  may  be  formed.  The  latter  is  found  as  NH4N03,  of  course, 
if  there  is  excess  nitric  acid  available  for  combination.  These 
results  agree  with  our  previous  experience  with  these  metals, 
since  we  have  already  found  in  the  study  of  the  electromotive 
series  that  zinc  gives  up  electrons  far  more  easily  than  does 
copper  (491). 

The  concentration  of  the  solution,  however,  as  well  as  the 
nature  of  the  reducing  agent,  determines  the  final  products  of 
these  reactions  by  reason  of  the  reversible  reaction 


Nitric  Acid  and  Oxides  of  Nitrogen  351 

We  find,  for  example,  that  NO  is  the  main  product  of  the  action 
of  copper  on  dilute  nitric  acid,  while  with  concentrated  nitric 
acid  NO2  is  chiefly  formed. 

558.  Reactions  of  Nitric  Acid  with  Metals. — All  the  common 
metals  except  platinum  and  gold  dissolve  in  nitric  acid  to  form 
nitrates,  though  the  latter  are  sometimes  much  hydrolyzed.     It 
might  be  thought  that  zinc  would  displace  hydrogen  from  nitric 
acid,  but  as  a  matter  of  fact  only  metals  in  the  electromotive 
series  down  to  and  including  aluminum  can  generate  hydrogen 
fast  enough  for  it  to  escape  the  oxidizing  action  of  nitric  acid. 
The  cases  of  copper  and  zinc  already  discussed  may  be  taken  as 
illustrative  of  these  reactions. 

559.  Reactions  of   Nitric   Acid  with  Non-Metals. — Finely 
divided  sulfur  is  oxidized  to  sulfuric  acid  by  hot  nitric  acid. 
Iodine  is  oxidized  to  iodic  acid,  HIO3.     Red  phosphorus  is 
oxidized  to  phosphoric  acid  (Care !  see  591).     Powdered  charcoal 
(carbon)  is  oxidized  to  carbon  dioxide.     In  all  of  these  actions 
brown  fumes  are  given  off,  showing  the  release  of  NO  or  NO2. 

560.  Nitric  Acid  as  a  Chemical  Solvent  for  Salts. — The  fact 
that  nitric  acid  is  an  oxidizing  agent  accounts  for  its  ability  to 
dissolve  many  salts  on  which  hydrochloric  acid  fails  to  act. 
Thus  the  insoluble  copper  sulfide  does  not  dissolve  appreciably  in 
HC1  (456)  but  does  dissolve  in  dilute  nitric  acid,  forming  copper 
nitrate  and  sulfur,  or  sulfuric  acid  if  the  action  with  the  acid  is 
continued.    So  also  mercurous  chloride,  insoluble  in  HC1,  is 
dissolved  easily  by  dilute  HNO3  to  form  mercuric  nitrate.     If 
neither  of  the  ions  of  a  salt  can  be  oxidized  nitric  acid  is  no  better 
chemical  solvent  than  HC1.    Thus  neither  acid  can  dissolve 
barium  sulfate  appreciably. 

561.  Balancing  of  Oxidation  and  Reduction  Equations. — 
Before  leaving  the  subject  of  oxidation  by  nitric  acid  it  is  worth 
while  to  take  up  a  systematic  method  for  balancing  oxidation 
and  reduction  equations,  to  be  used  when  the  proper  coefficients 
cannot  be  easily  arrived  at  by  inspection.    The  following  illus- 
tration of  a  useful  method  is  given  in  detail.    As  has  been  said, 
copper  reacts  with  concentrated  nitric  acid  to  give  copper  nitrate, 
nitrogen  tetroxide,  and  water.    We  may  first  write  down  the 


352  Introduction  to  General  Chemistry 

formulae  of  the  substances  taken  and  of  the  products  in  the  form 
of  an  equation  without  the  coefficients  : 


Inspection  of  the  formulae  shows  that  the  only  atoms  which 
change  their  valence  are  those  of  copper  and  those  of  nitrogen 
which  go  to  form  NO2  (not  those  which  go  to  form  nitrate)  .  The 
equation  for  the  exchange  of  charges  by  these  atoms  is  as  follows: 


The  equation  is  balanced  so  that  the  number  of  electrons  given 
up  by  the  copper  atoms  equal  those  taken  by  the  nitrogen.  We 
may  next  write  the  equation  to  represent  the  change  between  the 
molecules  which  contain  these  atoms: 

Cu+2HNO3->Cu+++2N02+H2O+O—  .  (i) 

On  the  right-hand  side  we  need  not  trouble  to  assign  positive 
ions  to  negative  ions  to  form  molecules  unless  so  doing  accounts 
for  some  of  the  actual  products.  NO2  and  H2O  are  examples. 
We  may  next  write  equations  to  show  the  transformation  of  the 
Cu++  and  O  into  the  products  in  which  they  finally  appear, 

Cu++  +  2HNO3->Cu(NO3)2+  2H+  .  (2) 

Gathering  up  the  O  and  H+  from  (i)  and  (2),  we  may  com- 
bine them  to  form  water, 

O—  +2H+-»HaO.  (3) 

Since  all  the  products  are  accounted  for,  we  may  now  secure  the 
final  equation  by  adding  equations  (i),  (2),  and  (3)  and  canceling 
the  terms  which  appear  on  both  sides: 

Cu+4HNO3->Cu(N03)2+  2NOa+  2H20. 

A  little  experience  will  show  that  there  is  a  great  similarity 
in  oxidation  and  reduction  equations,  so  that  the  detail  which 
we  have  written  out  so  laboriously  will  soon  become  so  familiar 
that  the  intermediate  equations  need  not  be  written  out.  How- 


Nitric  Acid  and  Oxides  of  Nitrogen  353 

ever,  until  this  facility  is  gained  it  is  better  to  write  out  each  step, 
as  was  done  above. 

One  more  example,  which  is  a  difficult  one,  will  now  be  taken 
up,  namely  the  reaction  between  zinc  and  nitric  acid  to  form 
zinc  nitrate,  ammonium  nitrate,  and  water.  We  shall  follow 
exactly  the  same  scheme  as  before.  The  first  step  is  to  arrange 
the  initial  materials  and  end  products  in  the  form  of  an  equation 
without  the  coefficients: 


Next  we  write  and  balance  an  equation  to  show  the  exchange  in 
charges  between  the  atoms  which  undergo  oxidation  and  reduc- 
tion: 


The  equation  for  the  reaction  of  the  molecules  which  contain 
these  atoms  is 

4Zn+HNO3->4Zn+++N  ---  +30  —  +H+.  (i) 

On  the  right-hand  side  of  the  equation  we  have  not  troubled  to 
assign  any  of  the  positive  ions  to  the  negative  ions,  since  in  so 
doing  we  would  not  account  for  any  of  the  actual  products. 
Next  we  must  account  for  the  actual  products  one  by  one. 
These  result  from  the  union  of  the  ions  represented  in  (i)  with 
the  ions  of  the  excess  acid  present: 

4Zn+++8HNO3^4Zn(NO3)2+8H+,  (2) 

N-  -+3H+->NH3,  (3) 

NH3+HNO3->NH4NO3.  (4) 

Gathering  up  the  excess  H+  and  O  from  equations  (i),  (2), 
(3),  and  (4)  we  may  write 

30—  +6H+-»3H20.  (5) 

All  products  are  now  accounted  for.  We  may  add  equations 
(i),  (2),  (3),  (4),  and  (5)  and  cancel  the  terms  which  appear  on 
both  sides  to  secure  the  final  equation 

ioHN03->4Zn(N03)2+NH4N03+3H2O. 


354  Introduction  to  General  Chemistry 

562.  Aqua  Regia. — As  has  already  been  stated,  gold  is  not 
attacked  by  nitric  acid;  but  this  royal  metal  dissolves  readily  in 
the  liquid  made  by  mixing  nitric  acid  with  three  times  its  volume 
of  concentrated  hydrochloric  acid.     Therefore  the  liquid  solvent 
for  the  royal  metal  was  called  aqua  regia  ("royal  water")  by  the 
alchemists  of  a  thousand  years  ago.     When  aqua  regia  is  gently 
warmed  the  liquid  becomes  yellow  and  gives  off  gases  as  the  result 
of  a  reaction  probably  best  represented  by  the  equation 

3HCl+HNO3->2H/)+NOCl+Cl2. 

The  substance  NOC1  is  called  nitrosyl  chloride.  The  action  of 
aqua  regia  on  gold  converts  the  latter  into  auric  chloride,  AuCl3> 
a  yellow  salt  easily  soluble  in  water.  Platinum,  which  like  gold 
is  also  insoluble  in  either  nitric  or  hydrochloric  acid  singly,  dis- 
solves in  aqua  regia  to  form  platinic  chloride,  PtCl4. 

563.  Nitrosyl  Chloride,  NOC1. — This  substance  is  a  brownish- 
colored  gas  of  very  disagreeable  odor.     It  is  used  extensively 
for  bleaching  flour.    Less  than  one  gram  of  NOC1  is  required  to 
bleach  a  barrel  of  flour.     Nitrosyl  chloride  is  readily  acted  on 

by  water,  thus : 

NOC1+H2O->HC1+HN02. 

The  nitrous  acid  formed  is  unstable  and  soon  decomposes: 
3HNO2-»HNO3+  2NO+H2O. 

564.  Nitrates. — Nearly  every  basic  ion  can  form  a  nitrate. 
The  nitrates  are  all  easily,  some  extremely,  soluble  in  water.    A 
few  nitrates  derived  from  weak  bases  are  hydrolyzed  by  water, 
forming  difficultly  soluble  basic  nitrates.    Thus  mercuric  nitrate 
and  water  give  Hg(N03)2'2HgO-H20,  called  basic  mercuric 
nitrate.    Bismuth  nitrate   and  water  give   the  basic  nitrate 
BiONO3,  called  bismuth  subnitrate,  used  extensively  in  medicine. 

The  nitrate  ion,  N03~,  is  colorless,  and  so  also  are  all  nitrates 
of  colorless  basic  ions. 

565.  Decomposition  of  Nitrates  by  Heat. — All  nitrates  are 
decomposed  when  they  are  heated  to  sufficiently  high  tempera- 
tures.   They  fall  into  three  classes  with  respect  to  their  behavior 
when  heated. 


Nitric  Acid  and  Oxides  of  Nitrogen  355 

1.  When  sodium  nitrate,  NaNO3,  is  heated,  it  first  melts 
without  decomposition  to  give  a  colorless  liquid.     At  a  higher 
temperature  the  liquid  appears  to  boil.     The  gas  given  off  is 
oxygen,  formed  as  follows: 

2NaNO3->2NaNO2-f-O2. 

The  other  product,  NaNO2,  is  sodium  nitrite,  the  formation  of 
which  by  the  action  of  NO2  on  NaOH  has  already  been  mentioned 
(553)-  Potassium  nitrate,  KNO3,  when  heated  gives  potassium 
nitrite,  KNO2,  and  oxygen. 

2.  The  effect  of  heat  on  lead  nitrate  results  in  the  following 
decomposition  : 

2Pb(N03)2->2PbO+4NO2+O2. 

This  reaction  is  typical  of  the  behavior  of  the  nitrates  of  most 
metals  excepting  those  of  the  alkali  group  (Na,  K,  etc.).  In 
some  cases  the  oxide  of  this  metal  is  also  decomposed  by  heat, 
so  that  the  free  metal  is  formed.  This  is  the  case  with  the 
nitrates  of  mercury  and  silver. 

3.  The  action  of  gentle  heat  on  ammonium  nitrate,  which 

takes  place  thus, 

NH4N03->N20+2H20, 

has  already  been  described  (556). 

566.  The  Equilibrium  between  Nitrogen  and  Oxygen.  —  We 
are  now  ready  to  take  up  in  detail  the  consideration  of  the  prob- 
lem of  the  manufacture  of  nitric  acid  from  atmospheric  nitrogen. 
We  shall  begin  by  the  study  of  the  reversible  reaction 


This  reaction  takes  place  so  slowly  below  800°  (a  bright-red  heat) 
that  ordinary  observation  would  lead  one  to  conclude  that 
nitrogen  and  oxygen  have  no  tendency  to  unite,  and  that  nitric 
oxide  has  no  tendency  to  dissociate  into  its  constituents.  At 
very  high  temperatures  each  reaction  takes  place  more  rapidly, 
so  that  at  the  very  high  temperature  of  3000°  equilibrium  is 
reached  in  a  small  fraction  of  a  second.  At  3000°  only  about 
5  per  cent  of  the  N2  and  O2  are  combined  as  NO  in  the  equilibrium 


356  Introduction  to  General  Chemistry 

mixture.  At  lower  temperatures  the  equilibrium  mixture  con- 
tains still  smaller  proportions  of  NO.  Table  XXIV  shows  the 
percentage  by  volume  of  NO  in  the  equilibrium  mixture  at  each 
of  several  temperatures,  and  also  the  length  of  time  required  to 
reach  practically  complete  equilibrium  (that  is,  within  one-tenth 
of  i  per  cent  complete).  Inspection  of  the  table  shows  that  the 
higher  the  temperature  the  greater  the  proportion  of  NO  formed 
and  the  quicker  the  attainment  of  equilibrium. 

TABLE  XXIV 


Temperature 

Per  Cent  of  NO 

Time 

o 
IxOO    

O.  I 

12  days 

2000° 

I    2 

30  sec 

2^00° 

2    6 

o  i  sec. 

3      0 

3OOO    

e     ? 

o  0004  sec. 

567.  The  Manufacture  of  Nitric  Acid  from  Air. — It  is  a 
simple  matter  to  make  nitric  acid  from  nitric  oxide,  NO.  We 
have  learned  that  NO  unites  readily  at  ordinary  temperatures 
with  oxygen  of  the  air  to  form  nitrogen  tetroxide, 

2NO->O2+2NO2. 

The  latter  gas  reacts  with  water  to  give  nitric  acid  and  nitric 

oxide, 

3NO2+H2O->2HNO3+NO; 

and  since  the  NO  by  direct  union  with  oxygen  of  the  air  passes 
readily  into  NO2,  the  whole  of  the  NO  can  finally  be  converted 
into  nitric  acid.  The  reactions  discussed  in  this  and  the  fore- 
going paragraphs  are  now  carried  out  on  a  very  large  scale  for 
the  manufacture  of  nitric  acid  from  atmospheric  nitrogen.  The 
difficult  part  of  the  process  is  the  preparation  of  NO.  We  see 
from  the  table  that  the  higher  the  temperature  to  which  the 
mixture  of  N2  and  O2  is  heated  the  greater  the  proportion  of 
NO  present  in  the  equilibrium  mixture.  Now  it  is  not  at  all 
difficult  to  heat  air,  which  is  simply  a  mixture  of  N2  and  O2,  to 
3000°  or  even  higher  by  means  of  an  electric  arc.  The  air  at 
this  temperature  then  contains  5  per  cent  or  more  of  NO,  which 


Nitric  Acid  and  Oxides  of  Nitrogen 


357 


is  a  very  satisfactory  proportion.  But  a  most  serious  difficulty 
is  now  encountered :  this  gas  mixture  is  at  a  dazzling  white  heat, 
3000°!  It  must  be  cooled  almost  to  ordinary  temperatures 
before  it  will  unite  with  more  oxygen  to  form  N02 ;  and  the  latter 
must  be  nearly  cold  before  it  is  combined  with  H2O  to  produce 
HN03.  Since  the  cool- 
ing cannot  be  accom- 
plished instantaneously, 
the  temperature  for  an 
appreciable  time  will  be 
between  2  500°  and  2000°. 
Therefore  the  reaction 


will  reverse  to  a  greater 

or  less  extent,  since  at 

2500°  only  2.6  per  cent 

of  NO  is  stable  and  at 

2000°  only  i .  2  per  cent. 

If    the    mixture   which 

contained  5  per  cent  of 

NO  at  3000°  remains  at 

2000°  for  30  seconds  only 

i .  2  per  cent  of  NO  will  FIG.  90 

be  left!    This  makes  it 

imperative  to  cool  the  gas  mixture  very  rapidly  to  a  temperature 

below  1000°,  where  the  speed  of  decomposition  of  NO  becomes 

negligible. 

One  of  the  most  successful  practical  methods  of  making  NO 
will  now  be  described. 

568.  The  Birkland  and  Eyde  Process. — This  process  was 
invented  in  1903  by  two  Norwegians,  Birkland  and  Eyde.  Air 
is  passed  through  an  immense  electric  arc  and  then  cooled  as 
rapidly  as  possible.  The  mixture  of  unchanged  air  and  nitric 
oxide  is  then  used  to  produce  nitric  acid  by  means  of  the  reactions 
already  discussed  (567).  The  Birkland-Eyde  electric  furnace 
is  shown  in  cross-section  in  Fig.  90.  An  arc  formed  at  A  between 


358 


Introduction  to  General  Chemistry 


the  electrodes  which  enter  from  the  front  and  back  is  spread  out 
by  the  action  of  a  powerful  magnet  N-S  into  a  great  electric 
flame  which  fills  the  space  B-B.  The  wires  of  the  magnet  are 
shown  at  M-M.  The  arc  is  inclosed  by  refractory  walls,  pierced 
by  numerous  holes.  (A  refractory  substance  is  one  which  is 
incombustible  and  may  be  heated  to  a  high  temperature.)  Air 
passes  through  the  furnace  in  the  manner  shown  by  the  arrows. 
The  temperature  of  the  perforated  walls  inclosing  the  arc  does 
not  exceed  800°,  owing  to  the  cooling  effect  of  the  inflowing  air. 
In  the  arc  the  gases  reach  a  temperature  of  3500°,  but  are  sub- 
sequently quickly  cooled  by  the  walls,  so  that  they  escape  at 


FIG.  91 

about  1000°.  The  NO  content  is  about  2  per  cent.  In  addition 
to  the  furnace  just  described,  several  others  accomplishing 
similar  results  are  in  practical  use. 

569.  Conversion  of  Nitric  Oxide  into  Nitric  Acid.  —  Figure  91 
illustrates  diagrammatically  a  typical  plant  for  the  manufacture 
of  nitric  acid  from  air. 

A  pump,  -A  (air  compressor),  delivers  air  to  the  electric 
furnace,  F,  from  which  the  escaping  mixture  of  air  and  NO  at 
1000°  passes  into  a  cooler,  C,  where  the  temperature  is  lowered 
to  500°.  The  gases  are  now  passed  under  steam  boilers,  B,  to 
generate  steam  for  use  in  the  plant,  and  are  thereby  still  further 
cooled.  The  reaction 


now  takes  place  at  about  50°  in  the  vessel,  D,  there  being  a  large 
excess  of  Oa  still  present.    The  absorption  tower,  T,  into  the 


Nitric  Acid  and  Oxides  of  Nitrogen  359 

base  of  which  the  gases  now  pass,  is  filled  with  pieces  of  quartz 
and  is  supplied  with  a  stream  of  water  at  the  top.  Here  the 

reaction 

3NO2+H2O->2HNO3-fNO 

takes  place.  Nitric  acid  is  drawn  off  at  the  base  of  the  tower, 
while  the  NO  and  air  pass  out  at  the  top.  The  escaping  NO  is 
not  lost,  since  it  can  combine  with  more  O2  to  form  N02,  which  is 
also  converted  into  nitric  acid,  so  that  finally  all  NO  is  changed 
into  HNO3.  Large  amounts  of  nitric  acid  are  now  manufactured 
in  this  way. 

570.  The  Manufacture  of  Nitric  Acid  from  Ammonia. — The 
conversion  of  NH3  into  HN03  was  mentioned  in  the  foregoing 
chapter.    When  a  mixture  of  one  volume  of  NH3  and  two 
volumes  of  oxygen  or  ten  volumes  of  air  is  passed  over  a  gauze 
of  fine  platinum  wire  at  about  700°  the  following  reaction  takes 
place : 

NH3+  202->HNO3-}-H2O. 

The  platinum  acts  catalytically,  and  the  speed  of  the  reaction  is 
so  great  that  contact  with  the  gauze  for  o. 01  second  is  sufficient. 
Considerable  heat  is  given  out,  so  that  by  regulation  of  the  rate 
of  flow  of  the  gases  the  most  favorable  temperature  is  maintained. 
Although  this  reaction  has  been  known  since  1830,  it  did  not 
assume  commercial  importance  until  after  cheap,  synthetic 
ammonia  became  available.  Since  the  beginning  of  the  war  a 
large  proportion  of  the  nitric  acid  required  by  Germany  has  been 
made*  from  NH3  by  this  process.  As  stated  in  the  foregoing 
chapter  (526),  an  immense  amount  of  nitric  acid  will  be  made 
at  the  new  American  works  at  Muscle  Shoals,  Alabama,  from 
NH3  made  by  the  cyanamid  process. 

571.  Uses  of  Nitric  Acid.    Explosives. — With  very  few  ex- 
ceptions all  practical  explosives  are  made  by  the  use  of  nitrates, 
or  nitric  acid.    The  oldest  explosive  is  gunpowder  (now  better 
known  as  black  powder).    This  is  a  mechanical  mixture  of 
charcoal,  sulfur,  and  potassium  nitrate.    Modern  explosives  are 
made  by  the  action  of  nitric  acid  on  various  substances,  such  as 
glycerine,  cotton,  phenol,  toluene,  etc.    The  chemistry  of  these 


360  Introduction  to  General  Chemistry 

nitrated  substances,  all  of  which  are  carbon  compounds,  will  be 
considered  in  chapter  xxvi.  Suffice  it  to  say  here  that  without 
nitric  acid  we  should  have  no  modern  explosives,  and  without 
modern  explosives  present  forms  of  warfare  would  not  exist. 
We  should  also  remember,  however,  that  the  arts  of  peace  are 
quite  as  dependent  upon  explosives  as  are  those  of  war.  Modern 
methods  of  mining  and  quarrying  would  be  impossible  without 
explosives;  and  without  the  aid  of  these  marvelous  chemical 
products  the  Panama  Canal  could  not  have  been  dug. 

Nitric  acid  is  an  indispensable  reagent  in  the  manufacture 
of  most  artificial  dyes.     It  also  has  many  other  important  uses. 

572.  Gunpowder,  Black  Powder. — The  history  of  the  dis- 
covery of  gunpowder,  usually  ascribed  to  Roger  Bacon,  is  rather 
obscure.     Gunpowder  seems  first  to  have  been  used  in  cannon 
toward  the  end  of  the  thirteenth  century.     This  explosive  is 
a  black,  granular  substance,  composed  of  about  74  per  cent 
potassium  nitrate,  16  per  cent  charcoal,  and  10  per  cent  sulfur. 
In  making  gunpowder  the  finely  powdered  components  are 
moistened  with  water,  and  the  mass  is  thoroughly  mixed.    The 
product  is  pressed  into  cakes;   these  are  dried  and  broken  into 
grains  of  suitable  size.     The  explosion  of  gunpowder  results  in 
the  change  of  KN03,  S,  and  C  into  C02,  S02,  N2,  and  K2S.    Of 
the  products  all  but  the  potassium  sulfide  are  gases.    The  very 
large  volume  of  gases  formed  in  a  fraction  of  a  second  is  the  cause 
of  the  result  which  we  call  an  explosion.     These  gases,  at  the 
instant  of  their  liberation,  are  at  a  very  high  temperature  and 
occupy  only  the  same  space  as  that  of  the  powder.     They*  there- 
fore exert  an  enormous  pressure  on  all  sides.     In  a  gun  the  pro- 
jectile is  shot  out,  while  the  gun  itself  recoils. 

573.  Ammonium    Nitrate    as    an    Explosive. — Ammonium 
nitrate,  NH4N03,  is  a  very  important  explosive.     We  have  seen 
that  when  gently  heated  it  decomposes,  thus : 

NH4NO3^N2O-h2H2O; 

but  when  heated  to  a  very  high  temperature  or  when  exploded 
by  a  powerful  detonator  the  decomposition  is  more  complete: 

2NH4N03->2N3+4H2O4-O2. 


Nitric  Acid  and  Oxides  of  Nitrogen  361 

At  the  high  temperature  of  the  explosion  H20  is  of  course  gaseous. 
A  detonator  is  a  substance  that  explodes  readily  and  produces  a 
sharp  shock  that  starts  the  explosion  of  a  large  quantity  of 
another  less  easily  exploded  substance.  Ammonium  nitrate  is 
not  easily  exploded.  It  may  therefore  safely  be  used  in  shells 
fired  from  cannon  without  risk  of  its  premature  explosion.  A 
time  or  impact  detonator  is  required  for  such  shells.  Enormous 
quantities  of  NH4N03  were  used  in  the  war. 

574.  The  Cycle  of  Nitrogen  in  Nature. — We  have  already 
learned  (511)  that  the  flesh  of  all  animals  is  made  up  of  com- 
pounds of  nitrogen  (with  carbon,  hydrogen,  and  oxygen  chiefly). 
Animals  cannot  derive  their  required  nitrogen  from  the  air  or 
from  any  of  the  simple  nitrogen  compounds  so  far  studied. 
Much  more  complex  nitrogen  compounds  (proteins)  are  required 
for  the  food  of  animals  (including  man).  These  nitrogenous 
animal  foods  are  obtained  from  the  flesh  of  other  animals  and 
from  plants,  particularly  cereals.  Plants  have  the  power  of 
building  up  proteins  from  simple  nitrogen  compounds,  such  as 
ammonium  salts  and  nitrates,  which  are  present  in  the  soil.  A 
small  group  of  plants  (especially  the  legumes:  clover,  alfalfa, 
etc.)  can,  by  virtue  of  bacteria  that  infest  their  roots,  take  up 
free  nitrogen  and  convert  it  into  proteins.  All  other  plants 
thrive  only  on  soils  containing  compounds  of  nitrogen.  Soils 
deficient  in  combined  nitrogen  are  greatly  improved  in  fertility 
by  the  application  of  ammonium  salts  or  of  nitrates  as  fertilizers. 
In  normal  times  about  three-fourths  of  all  ammonium  salts  and 
nitrates  used  in  America  are  employed  as  fertilizers.  A  large 
part  of  the  nitrogen  of  animal  food  is  eliminated  in  the  form  of  a 
rather  simple  substance,  urea,  CO(NH2)2.  This  substance 
changes  slowly  into  C02  and  NH3  by  the  action  of  water, 

CO(NH2)2+H20->C02+  2NH3. 

Urea  itself,  as  well  as  NH3,  can  serve  as  plant  food. 

The  decay  of  animal  remains  and  refuse  results  in  the  change 
of  their  proteins  into  simpler  substances.  Ammonium  salts  are 
first  formed  and  later  are  oxidized  to  nitrates.  Thus  nitrogen 
in  nature  passes  through  an  endless  cycle  of  changes. 


362  Introduction  to  General  Chemistry 

575.  The  Problem  of  the  Fixation  of  Nitrogen. — In  this 
chapter  and  the  foregoing  chapters  we  have  tried  to  point  out 
the  practical  importance  of  the  compounds  of  nitrogen.  The 
abundance  of  food  is  determined  by  the  fertility  of  the  soil. 
Fertility  may  be  conserved  and  increased  by  means  of  nitrogenous 
fertilizers.  Aside  from  animal  refuse  and  manures  and  the 
ammonia  obtained  as  a  by-product  of  the  manufacture  of  gas 
and  coke,  the  great  deposits  of  sodium  nitrate  in  Chile  were 
until  recently  the  only  important  source  of  nitrogenous  fer- 
tilizers. These  deposits  are  by  no  means  inexhaustible.  They 
may  last  one  hundred  years  at  the  present  rate  of  consumption. 
Fortunately,  within  the  last  twenty-five  years  several  methods 
have  been  developed  for  making  nitrogen  compounds  directly 
from  the  free  nitrogen  of  the  air.  All  such  methods  are  now 
referred  to  as  processes  for  the  fixation  of  nitrogen.  We  have 
at  present  three  successful  methods,  each  of  which  is  in  extensive 
use.  These  are:  (i)  the  making  of  nitric  acid  from  air  (566), 
(2)  the  making  of  cyanamid  from  calcium  carbide  and  atmos- 
pheric nitrogen  (526),  and  (3)  the  synthesis  of  ammonia  from 
hydrogen  and  atmospheric  nitrogen  (520).  We  have  also  seen 
that  ammonia  is  readily  made  from  cyanamid,  and  that  nitric 
acid  is  easily  produced  by  the  oxidation  of  ammonia  by  air. 
Since  the  nitrogen  contained  in  the  air  over  each  square  mile  of 
the  earth's  surface  amounts  to  nearly  20,000,000  tons,  the  supply 
may  be  said  to  be  inexhaustible. 

The  final  question  to  be  settled  before  deciding  upon  the 
practicability  of  any  technical  process  is  whether  it  is  eco- 
nomically sound;  in  other  words,  will  it  pay?  With  regard  to 
the  three  general  methods  of  fixing  nitrogen  above  mentioned, 
it  may  be  said  that  the  direct  manufacture  of  nitric  acid  from 
air  requires  exceedingly  cheap  electric  power.  The  other  two 
general  methods  can  be  used  where  power  is  more  expensive. 
At  present  all  three  processes  are  profitable.  Only  time  will 
decide  which  will  prove  most  economical.  In  any  case  we  may 
feel  satisfied  that  the  problem  of  the  fixation  of  nitrogen  has  been 
solved. 


CHAPTER  XXIII 
PHOSPHORUS 

576.  Review. — We  have  already  had  some  experience  with 
the  chemistry  of  phosphorus  and  its  compounds.     The  burning 
of  the  free  element  was  used  to  show  the  presence  of  oxygen  in 
air  (10).     In  chapter  ix  phosphoric  acid  and  some  of  its  salts 
were  described  (158,  159).     The  importance  of  phosphates  for 
the  life  of  plants  and  animals  was  pointed  out  (160),  and  their 
widespread  distribution  in  nature  was  mentioned.     One  method 
of  preparing  free  phosphoric  acid  was  also  given.     In  chapter  xii 
(247)  the  chemistry  of  phosphorus  was  continued  in  the  descrip- 
tion of  the  union  of  phosphorus  with  the  halogens  to  form 
trivalent  and  pentavalent  phosphorous  compounds,  and  the 
reactions  of  these  compounds  with  water  to  set  free  the  hydrogen 
halides  and  an  oxygen  acid  of  phosphorus,  as  illustrated  by  the 
following  typical  equations: 

PCl3+3H20->H3P03-HHei, 
PBrs+4H2O->H3PO4+  5HBr. 

In  chapter  xxii  the  powerful  dehydrating  property  of  the  pen- 
toxide  of  phosphorus  was  mentioned  in  the  formation  of  the 
anhydride  of  nitric  acid: 

2HNO3+P2O5-»2HPO3+N20S.  (555) 

We  shall  now  take  up  a  systematic  study  of  the  element 
phosphorus  and  its  compounds. 

577.  The   Discovery  of  the  Element. — The  free  element, 
which  does  not  occur  in  nature,  was  discovered  by  Brandt,  of 
Hamburg,  sometime  before  1669.      In  searching  for  a  suitable 
solvent  for  turning  silver  into  gold  he  was  led  to  ignite  a  mixture 
of  dried  urine  and  charcoal  in  a  clay  retort.    Since  the  urine 
contained  phosphates  he  was  really  carrying  out  one  of  the 

363 


364  Introduction  to  General  Chemistry 

methods  now  in  use  for  preparing  free  phosphorus.  When  we 
have  studied  some  of  the  peculiarities  of  the  free  element  we  shall 
realize  what  a  stir  this  new  substance  must  have  made  in  the 
thinking  world  of  that  day.  Historians  record  with  unction 
that  Kraft,  who  purchased  the  secret  of  its  preparation  from 
Brandt,  exhibited  it  before  many  of  the  crowned  heads  of  Europe. 

578.  The  Physical  Properties  of  White  Phosphorus. — At 
ordinary  temperatures  the  free  element  in  pure  form  is  a  white, 
waxy,  translucent  solid  of  specific  gravity  1.8.     It  does  not  con- 
duct electricity.     It  is  usually  kept  under  water,  since  when  it 
is  exposed  to  air  it  soon  takes  fire.    At  ordinary  temperatures 
the  solid  is  slightly  volatile.     Calculations  made  from  the  density 
of  the  vapor  show  that  it  has  a  formula  of  P4  at  temperatures  up 
to  1,500°.     It  melts  at  45°  and  boils  at  287°.    It  is  not  soluble  in 
water  but  is  readily  soluble  in  many  solvents,  such  as  carbon 
disulfide,  turpentine,  olive  oil,  and  many  of  its  own  compounds. 
The  so-called  yellow  phosphorus  of  commerce  is  white  phos- 
phorus made  yellow  by  the  presence  of  impurities. 

579.  The  Conversion  of  White  Phosphorus  into  Red  Phos- 
phorus.— When  white  phosphorus  is  exposed  to  light  it  darkens 
in  color,  owing  to  the  formation  of  a  new  form  called  from  its 
color  red  phosphorus.     Samples  of  white  phosphorus  which  have 
been  kept  for  some  time  show  a  coating  of  this  substance.    At 
higher  temperatures  the  change  from  white  to  red  goes  on  more 
rapidly.     It  can  be  hastened  also  by  the  presence  of  a  trace 
of  iodine,  which  acts  as  a  catalyzer. 

580.  The  Manufacture  of  Phosphorus. — Though  some  phos- 
phorus is  manufactured  from  bone  ash,  most  of  it  is  prepared 
from  phosphate  minerals.    The  principal  materials  are  phos- 
phate rock  and  apatite.    The  first  varies  in  composition  from 
Ca3(PO4)2  to  the  composition  of  apatite,   Ca3(P04)2CaF2,   or 
Ca3(P04)2CaQ2. 

The  most  widely  used  method  of  making  phosphorus  con- 
sists in  heating  a  mixture  of  sand,  phosphate  rock,  and  charcoal 
to  high  temperatures  in  an  electric  furnace,  so  designed  that  the 
phosphorous  vapor  distils  off  and  condenses  under  water.  The 
residual  materials  of  the  reaction  are  withdrawn  from  time  to 


Phosphorus  365 

time,  and  new  mixtures  are  put  in  in  a  continuous  process.     The 
equation  of  the  reaction  follows: 

2Ca3(PO4)2+6SiOa+ioC^6CaSiO3+ioCO+P4. 

Red  phosphorus  is  made  on  a  commercial  scale  by  heating 
yellow  phosphorus  in  the  absence  of  air. 

581.  The   Properties  of  Red   Phosphorus.— Although   red 
phosphorus  is  often  given  the  title  amorphous  phosphorus,  which 
means  that  it  is  without  crystalline  form,  it  is  really  a  crystalline 
powder.    Its  specific  gravity  is  2.1.     Commercial  red  phos- 
phorus often  contains  among  other  impurities  small  traces  of 
yellow  phosphorus  and  some  phosphoric  acid.     It  is  on  this 
account  sticky  and  hygroscopic  (water-absorbing),  although  if 
free  from  these  impurities  it  is  a  dry  powder  quite  insoluble  in 
water  and  also  in  most  other  solvents.     It  has  far  greater  stability 
than  white  phosphorus,  as  is  evidenced  by  the  fact  that  it  must 
be  heated  to  240°  before  it  will  take  fire  in  air,  and  that  it  may 
be  kept  in  contact  with  the  latter  at  ordinary  temperatures  for 
years  without  alteration.     If  pure  it  is  non-poisonous.     It  is 
for  these  reasons  to  be  preferred  to  white  phosphorus  in  all 
laboratory  experiments  in  which  the  substitution  can  be  made. 
The  chemistry  of  red  phosphorus  parallels  that  of  yellow  phos- 
phorus, but  in  every  case  the  reactions  are  less  violent. 

582.  Allo tropic  Forms. — White  and  red  phosphorus  are  called 
allotropic  forms  of  phosphorus.    The  noun  allotropy  translated 
means  simply  "other  modes."     We  shall  find  many  examples  of 
elements  which  exist  in  more  than  one  form.    We  have  already 
met  ordinary  oxygen  and  ozone  (316),  allotropic  forms  of  the 
same  element.     We  found  that  these  two  forms  differed  from 
each  other  in  the  number  of  atoms  and  in  the  energy  content  of 
the  molecule.     The  same  causes  undoubtedly  account  for  the 
differences  in  the  two  forms  of  phosphorus. 

583.  The  Slow  Oxidation  of  Phosphorus. — If  white  phos- 
phorus is  exposed  to  moist  air  the  heat  given  out  by  its  oxidation 
will  slowly  raise  its  temperature  to  the  melting-point,  and  it  then 
takes  fire.     It  is  this  fact  that  makes  the  storage  of  white 
phosphorus  under  water  necessary. 


366  Introduction  to  General  Chemistry 

If  a  stick  of  moist  phosphorus  is  viewed  in  the  dark  it  is  seen 
to  give  out  a  soft,  yellow  light.  It  is  from  this  property  that 
the  element  received  its  name,  which  means  "bearer  of  light." 
We  can  imagine  what  an  interest  the  discovery  of  this  substance 
aroused  in  a  superstitious  age !  The  appearance  of  this  light  is 
apparently  due  to  the  slow  oxidation  which  is  in  progress;  for  if 
conditions  are  such  that  oxidation  cannot  occur,  no  light  is 
given  out  by  the  phosphorus.  The  light  is  simply  a  part  of  the 
energy  of  the  reaction,  which  appears  in  this  unusual  form 
instead  of  as  heat. 

584.  Phosphorescence. — When  light  is  given  off  by  bodies  at 
ordinary  temperatures  they  are  said  to  phosphoresce,  and  the 
process  is  called  phosphorescence.    But  the  same  phenomenon 
is  shown  by  substances  other  than  phosphorus.    Moist,  decaying 
wood  often  glows  in  the  dark  as  the  result  of  the  liberation  of 
energy  in  the  form  of  light.    The  so-called  fox  fire  seen  in 
forests  after  night  is  an  example.    We  have  another  case  in  the 
process  by  which  the  firefly  gives  out  light  from  parts  of  the 
surface  of  its  body.    All  our  processes  of  making  light  are  very 
wasteful,  since  high  temperatures  must  be  maintained,  and  so 
most  of  the  energy  used  must  go  into  heat,  and  only  a  small 
portion  is  changed  into  light.    Hence  the  process  by  which  the 
firefly  gives  out  light  at  the  temperature  of  its  body  is  indeed 
marvelous,  and  its  imitation  has  been  the  despair  of  scientists 
for  generations. 

585.  Danger  in  the  Use  of  White  Phosphorus. — In  conclusion 
it  should  be  especially  emphasized  that  white  phosphorus  is  a 
very  dangerous  reagent.    It  should  never  be  touched  by  the  hands 
directly,  since  it  sticks  to  the  skin  and  the  heat  of  the  hand  is 
sufficient  to  ignite  it. 

586.  Matches. — On  account  of  the  serious  harm  to  workers 
with  yellow  phosphorus,  by  reason  of  poisonous  vapor,  the  mak- 
ing of  yellow-phosphorus  matches  has  been  made  illegal  in  all 
the  great  countries  of  the  world.    These  matches  were  at  one 
time  very  popular  because  of  their  easy  ignition.    Matches  are 
now  made  with  red  phosphorus,  phosphorus  sulfide,  P4S3,  or 
other  non-poisonous  compounds  of  phosphorus.    The  ordinary 


Phosphorus  367 

match  is  a  stick  of  non-resinous  wood  dipped  first  in  paraffin 
and  then  in  a  preparation  to  make  the  ignition  head.  The  latter 
is  usually  a  mixture  of  a  good  oxidizing  material  like  lead  dioxide 
or  potassium  chlorate  and  a  reducing  agent  such  as  red  phos- 
phorus or  phosphorus  sulfide.  With  these  is  mixed  some 
powdered  glass  to  increase  the  friction  when  the  match  is  struck. 
The  material  of  the  head  is  kept  in  place  by  glue  or  varnish. 
Friction  starts  the  action  between  the  oxidizer  and  the  reducer, 
and  the  heat  of  this  reaction  kindles  the  paraffin,  which  in  turn 
kindles  the  wood.  The  heads  of  safety  matches  are  made  of  less 
easily  reacting  mixtures,  usually  sulfur  and  potassium  chlorate. 
They  are  ignited  by  striking  on  a  prepared  surface  coated 
with  red  phosphorus,  antimony  trisulfide,  and  glue.  Match 
sticks  are  often  treated  with  phosphoric  acid  or  sodium  phosphate 
to  prevent  "afterglow,"  the  formation  of  glowing  coals  of  char- 
coal after  the  flame  of  the  match  has  been  blown  out. 

587.  Smoke   Screens. — A   new    use    for    phosphorus    was 
developed  during  the  war.     It  was  discovered  that  the  most 
effective  smoke  screens  for  concealing  troop  movements  were 
produced  by  means  of  burning  phosphorus.    The  explosion  of 
shells  filled  with  phosphorus  ignited  the  latter,  which  burned  to 
give  dense  white  clouds  of  the  pentoxide,  P2OS.    Large  quantities 
of  phosphorus  were  used  for  this  purpose. 

588.  Compounds  of  Phosphorus. — Phosphorus  forms  com- 
pounds with  many  elements.    With  hydrogen  it  unites  to  form 
several  hydrides.   Of  these,  phosphine,  PH3,  has  been  of  particu- 
lar interest,  since  it  apparently  resembles  ammonia.     It  is  a 
poisonous,  ill-smelling  gas,  little  soluble  in  water,  but  like  am- 
monia it  unites  with  acids  to  form  salts,  though  far  less  readily. 
For  instance,  phosphine  and  hydrogen  chloride  unite  at  low 
temperature  to  form  a  cloud  of  solid  phosphonium  chloride: 

PH3+HC1^PH4C1. 

Some  of  the  halogen  compounds  of  phosphorus  have  been 
already  studied.  These  are  very  important  substances  used  to 
synthesize  many  valuable  halogen  compounds.  The  formation 
of  hydrogen  bromide  or  hydrogen  iodide  by  the  action  of  water 


368  Introduction  to  General  Chemistry 

on  the  tribromide  or  the  tri-iodide  is  typical  of  these  important 
reactions  and  should  be  reviewed  (256,  264). 

With  oxygen,  phosphorus  unites  to  form  P203,  P204,  and 
P20S,  respectively.  The  first  two  are  formed  when  phosphorus 
burns  in  a  poor  supply  of  oxygen;  the  last  is  formed  when  the 
burning  takes  place  in  an  excess  of  oxygen.  All  three  are  white 
hygroscopic  solids. 

Phosphorus  trioxide  unites  with  cold  water  to  form  phos- 

phorous acid: 

3H20+P203^2H3P03. 

The  latter  is  a  very  unstable  acid.  When  the  solution  is 
heated  some  of  the  molecules  of  the  acid  are  oxidized  to  phos- 
phoric acid  at  the  expense  of  others,  which  are  reduced,  forming 
phosphine. 

Phosphorus  pentoxide  unites  with  water  with  great  vigor, 
forming  first  metaphosphoric  acid,  HPO3.  The  stability  of  this 
acid  and  the  completeness  of  the  reaction  make  phosphorus 
pentoxide  our  most  effective  drying  agent.  The  last  traces  of 
moisture  are  taken  from  air  exposed  to  this  oxide. 

589.  The  Phosphoric  Acids.  —  When  phosphorus  pentoxide 
is  thrown  into  water  spattering  is  likely  to  occur  because  of  the 
heat  given  out  in  the  formation  of  metaphosphoric  acid: 

H20+P205^2HP03. 

The  excess  water  may  be  evaporated,  leaving  the  pure  acid,  a 
colorless,  glassy  solid.  It  changes  to  orthophosphoric  acid, 
H3P04,  if  it  is  kept  a  long  time  in  water  solution.  The  change 
may  be  carried  out  in  the  course  of  an  hour  if  the  metaphos- 
phoric acid  solution  is  mixed  with  a  strong  acid  and  heated  on 
the  water  bath: 


The  action  is  hastened  by  the  catalytic  action  of  the  hydrogen 
ion  of  the  strong  acid. 

590.  The  Properties  of  Orthophosphoric  Acid.  —  Orthophos- 
phoric acid,  usually  called  simply  phosphoric  acid,  forms  color- 
less, odorless  crystals,  which  are  deliquescent  and  very  soluble 


Phosphorus  369 

in  water.  The  acid  is  usually  sold  in  the  form  of  a  concentrated 
solution  of  sirupy  consistency.  Its  dilute  solution  has  an  agree- 
able sour  taste  and  is  not  poisonous.  This  acid  cannot  be 
oxidized  and  is  very  stable  toward  most  reducing  agents.  It  is 
a  moderately  strong  acid  (409);  but  since  it  is  tribasic  (159)  we 
shall  consider,  in  some  detail,  in  a  separate  section,  its  interesting 
mode  of  ionization. 

Orthophosphoric  acid  is  not  volatile.  For  this  reason  it 
reacts  at  elevated  temperatures  with  salts  of  volatile  acids  like 
chlorides,  bromides,  and  nitrates  to  form  phosphates  and  to  set 
free  the  volatile  acids.  For  example, 

H3PO4+NaBr->NaH2PO4+HBr. 

When  phosphoric  acid  is  heated  to  2io°-2i5°  in  an  open  vessel 
for  about  an  hour  it  is  converted,  by  slow  loss  of  water,  into 
pyrophosphoric  acid,  H4P2O7: 

2H3P04->H4P2O7+H2O. 

If  the  heating  is  continued  longer  and  the  temperature  is  raised 
to  400°  or  higher,  the  product  slowly  loses  more  water,  giving 
finally  metaphosphoric  acid,  HP03: 

H4P207->2HP03+H2O. 

Metaphosphoric  acid  volatilizes  slowly  at  a  red  heat  (7oo°-8oo°) . 

591.  The  Preparation  of  Phosphoric  Acid. — Phosphoric  acid 
is  easily  made  in  the  laboratory  by  boiling  red  phosphorus  with 
concentrated  nitric  acid.  (Caution!  White  phosphorus  must 
not  be  used.)  The  ensuing  reaction  produces  red  fumes  of 
nitrogen  tetroxide,  NO2,  and  gives  finally  a  colorless,  sirupy 
solution,  from  which  the  excess  of  HNO3  can  be  removed  by 
heating,  since  HNO3  is  easily  volatile,  while  H3PO4  is  not. 
Incomplete  oxidation  gives  more  or  less  phosphorous  acid, 
H3PO3.  The  latter  is  readily  oxidized  to  H3PO4  by  sufficient 
boiling  nitric  acid. 

Phosphoric  acid  is  made  commercially  by  the  double  decom- 
position of  calcium  phosphate  and  sulfuric  acid  (158)  and  sub- 
sequent purification  of  the  product.  The  crude  acid  is  very 


370  Introduction  to  General  Chemistry 

likely  to  contain  magnesium  and  calcium  salts  and  also  some 
arsenic  acid,  H3As04,  the  first  coming  from  the  rock  used  and 
the  second  mostly  from  the  acid  employed,  though  sometimes 
from  the  rock  as  well. 

592.  lonization  of  Phosphoric  Acid.  —  According  to  the  ionic 
theory  phosphoric  acid  dissociates  as  follows: 

H3P04^H++H2P04-, 
H2P04-±5H++HPO4—  , 
HP04—  ^ 


In  a  water  solution  equilibrium  is  established  between  the  three 
pairs  of  opposed  reactions,  and  each  equilibrium  is  dependent 
on  the  other  two.  In  all  cases  of  this  kind  the  first  dissociation 
goes  on  to  a  greater  degree  than  the  second,  the  second  to  a 
greater  degree  than  the  third,  etc.  In  a  solution  of  phosphoric 
acid  containing  one-tenth  formula  weight  per  liter  at  18°  the 
total  concentration  of  hydrogen  ion  is  0.0282  and  that  of  the 
primary  phosphate  ion,  H2P04~",  virtually  the  same;  the  con- 
centration of  the  secondary  phosphate  ion,  H2P04  ,  is 
calculated  to  be  0.0000002,  while  that  of  the  phosphate  ion, 
PO4  ,  is  very  much  smaller. 

It  is  important  that  we  understand  that,  inasmuch  as  H2P04"~ 
and  HPO4  both  dissociate  to  give  hydrogen  ion,  they  are  acids, 
just  as  is  phosphoric  acid  itself.  Since,  however,  their  degrees  of 
dissociation  are  far  less  than  that  of  phosphoric  acid,  even  at 
greater  dilutions  they  are  far  weaker  acids.  Obviously  HP04  — 
is  much  weaker  than  H2P04~.  These  acids  are  furnished  by 
the  soluble  and  highly  ionized  salts  NaH2PO4  and  Na2HPO4  as 
a  result  of  their  ionization: 

NaH2PO4^Na++H2PO4~, 
4^  2Na+-hHPO4~  -  . 


593.  Phosphate  Baking  Powder.—  The  ion  H2P04~  is  a  weak 
acid  of  about  the  same  strength  as  carbonic  acid,  but  it  has  an 
advantage  over  the  latter  in  being  more  stable  in  water  solution. 
As  a  result,  when  water  is  added  to  a  mixture  of  sodium  hydrogen 
carbonate,  NaHC03,  and  sodium  acid  phosphate,  NaH2P04,  we 


Phosphorus  371 

get,  as  we  expect,  an  effervescence  due  to  the  escape  of  carbon 
dioxide  from  the  solution  of  carbonic  acid  formed.  Thanks  to 
the  escape  of  the  gas  the  action  becomes  complete: 

NaHCO3+NaH2PO4±5Na2HPO4+H2O+CO2. 

This  is  the  reaction  made  use  of  in  phosphate  baking  powders, 
The  essential  ingredients  are  sodium  bicarbonate  and  an  acid 
phosphate  which  may  be  the  sodium,  calcium,  or  magnesium 
salt,  or  mixtures  of  them,  together  with  some  inert  material, 
such  as  starch,  which  serves  as  a  dilutent.  When  water  is  added 
to  the  mixture  of  the  dry  salts  and  flour  the  tiny  bubbles  of 
carbon  dioxide  given  off  make  the  dough  spongy. 

594.  The  Phosphates.  The  Sodium  Salts. — Primary  sodium 
phosphate,  or  sodium  acid  phosphate,  NaH2P04,  is  prepared  in 
solution  by  adding  one  molecular  weight  of  sodium  hydroxide 
for  every  molecular  weight  of  phosphoric  acid : 

H3P04+NaOH->NaH2P04+H2O. 

That  NaH2P04  alone  forms,  and  not  a  mixture  of  the  three 
possible  phosphates,  we  can  understand  if  we  remember  that  the 
secondary  and  tertiary  salts  are  salts  of  weaker  acids  than 
H3P04;  and  so  if  momentarily  formed  they  would  undergo 
double  decomposition  with  unused  phosphoric  acid  to  form 
NaH2P04.  A  solution  of  NaH2PO4  is  acid  to  litmus,  as  we  would 
expect  from  the  presence  of  the  acid  H2P04~. 

Common  sodium  phosphate,  Na2HP04,  the  secondary  salt, 
is  prepared  by  adding  two  molecular  weights  of  sodium  hydroxide 
for  every  molecular  weight  of  phosphoric  acid  in  solution.  Solu- 
tions of  secondary  sodium  phosphate  are  alkaline  to  litmus, 
owing  to  a  slight  hydrolysis  (436). 

HP04  is  so  weak  an  acid  that  when  three  molecules  of 
sodium  hydroxide  are  furnished  for  every  molecule  of  phosphoric 
acid  in  solution  the  action  is  far  from  complete,  and  we  have  only 
a  small  part  of  the  possible  Na3P04  formed  when  equilibrium  is 
established.  However,  we  can  force  the  reaction  to  yield  the 
desired  salt  by  evaporating  off  the  water,  or  better  by  furnishing 


372  Introduction  to  General  Chemistry 

a  very  large  excess  of  sodium  hydroxide  and  then  evaporating 
the  water.  In  either  case  the  tertiary  or  normal  sodium  phos- 
phate will  separate  from  solution.  As  we  would  expect,  if  we 
redissolve  this  salt  we  get  back  the  same  solution  made  originally, 
and  we  find  that  as  a  result  of  the  hydrolysis  involved  the  solution 
is  strongly  alkaline. 

As  solids,  all  three  salts  are  white  crystalline  substances  and 
are  hydra  ted  as  indicated  by  the  following  formula: 

NaH2PO4-4H20;  Na2HP04-i2H2O;  Na3PO4-i2H2O. 

When  heated  the  primary  and  secondary  salts  both  give  off 
water  and  form  sodium  metaphosphate  and  pyrophosphate 
respectively.  The  tertiary  phosphate  remains  unchanged: 

NaH2PO4->NaPO3+H20, 
2Na2HPO4->Na4P2O7+H20. 

595.  The  Titration  of  Phosphoric  Acid,  Normal  Solution.— 

If  phenolphthalein  is  used  as  the  indicator  and  phosphoric  acid 
is  titrated  with  sodium  hydroxide  (137),  the  color  change  of  the 
indicator  will  be  observed  when  two  formula  weights  of  sodium 
hydroxide  have  been  added  for  every  formula  weight  of  the  acid. 
This  is  the  point  at  which  the  acid  has  been  changed  to  the 
secondary  salt,  Na2HPO4.  Therefore  a  liter  of  a  solution  of 
phosphoric  acid  which  contains  one-half  formula  weight  of  the 
acid  will  neutralize  one  liter  of  a  normal  sodium  hydroxide  solu- 
tion and  has  the  same  neutralizing  power  as  have  solutions  of 
hydrochloric  acid  or  nitric  acid  which  contain  one  formula 
weight  of  the  acid  per  liter  respectively,  or  of  sulfuric  acid  which 
contain  one-half  formula  weight  of  the  acid  per  liter.  A  half 
molar  solution  of  phosphoric  acid  is  therefore  a  normal  solution 
of  the  latter. 

596.  The  Bead  Test. — Sodium  metaphosphate  is  much  used 
in  analytical  chemistry  for  the  so-called  metaphosphate  bead 
tests.    The  bead  tests  are  carried  out  by  first  fusing  micro- 
cosmic  salt,  sodium  ammonium  hydrogen  phosphate,  on  the 
end  of  a  platinum  wire.     This  salt,  which  is  used  because  it 


Phosphorus  373 

is  easily  prepared  in  pure  form,  decomposes,  giving  sodium 
metaphosphate  as  follows: 

NaNH4HPO4->NaPO3-f-H20+NH3. 

Then  a  tiny  speck  of  the  material  to  be  tested  is  picked  up  on  the 
bead  and  fused  into  it.  Salts  of  metals  fuse  into  the  bead, 
forming  phosphates,  which  often  are  characteristically  colored 
and  thus  indicate  the  metal  present  in  the  original  substance. 
Presumably  the  salt  of  the  metal  first  forms  an  oxide  and  then 
unites  with  the  metaphosphate  to  form  an  orthophosphate,  a 
reaction  which  is  really  a  reverse  of  the  formation  of  the  meta- 
phosphate from  the  orthophosphate  described  above.  For 
instance,  copper  oxide  would  be  dissolved,  forming  sodium 
copper  phosphate,  thus: 

CuO+NaPO3->NaCuP04. 

597.  Qualitative    Tests    for    Phosphates. — Phosphates    are 
usually  tested  for  by  the  formation  of  yellow  ammonium  phos- 
phomolybdate,  (NH4)3PO4iiMo036H20,  which  is  insoluble  in 
nitric  acid  solution.    As  the  formula  indicates,  the  successful 
operation  of  this  test  necessitates  an  enormous  excess  of  the 
reagent  ammonium  molybdate. 

Owing  to  the  fact  that  solutions  of  the  metaphosphates  and 
pyrophosphates  are  stable  and  only  form  the  orthophosphate 
after  the  lapse  of  considerable  time,  unless  they  are  heated  with 
strong  acids,  it  is  desirable  to  be  able  to  distinguish  between 
the  three  kinds  of  phosphates.  This  is  done  on  the  basis  of  the 
fact  that  silver  orthophosphate,  Ag3PO4,  is  insoluble  in  water 
and  yellow,  whereas  the  other  silver  phosphates,  AgP03  and 
Ag4P207,  though  also  insoluble,  are  white.  The  metaphosphate 
is  distinguished  from  the  pyrophosphate  by  the  fact  that  the 
free  acid  which  can  be  made  from  the  salt  will  coagulate  albumen, 
commonly  known  in  the  form  of  white  of  egg,  while  pyrophos- 
phoric  acid  will  not.  As  a  matter  of  fact  this  test  distinguishes 
the  metaphosphoric  acid  from  the  orthophosphoric  acid  as  well. 

598.  Use  and  Production  of  Phosphates. — Phosphates  and 
phosphoric  acid  are  used  extensively  in  medicine,  but  the  great 


374  Introduction  to  General  Chemistry 

bulk  of  the  phosphate  production  supplies  the  match  industry 
and  the  still  greater  demands  of  agriculture.  The  meeting  of 
the  latter  need  is  of  incalculable  importance,  since  the  phos- 
phate taken  from  the  land  by  the  crops  must  be  put  back  or  the 
soil  will  become  sterile.  The  natural  phosphates  are  mined  on 
an  enormous  scale  either  for  direct  use  as  fertilizer  or  for  the 
production  of  superphosphate  (160),  to  be  used  for  the  same 
purpose. 


CHAPTER  XXIV 
SULFUR  AND  ITS  COMPOUNDS 

599.  Introduction. — Many  isolated  cases  of  the  use  of  sulfur 
compounds    have    already   been   encountered.     Some   of    the 
properties  of  sulfuric  acid  were  given  (93),  and  in  subsequent 
work  its  use  was  described  in  the  preparation  of  hydrochloric 
(103),  nitric  (104),  and  phosphoric  (158)  acids,  and  also  in  the 
making  of  many  sulfates  and  acid  salts.    Hydrogen  sulfide  was 
used  as  a  reducing  agent  (339),  the  most  important  instance 
being  the  preparation  of  hydrogen  iodide  by  the  action  of 
aqueous  hydrogen  sulfide  upon  iodine.     Sulfurous  acid  was  also 
shown  to  be  a  reducing  agent  because  of  its  easy  oxidation  to 
sulfuric  acid  (340) .    As  a  matter  of  fact,  so  commonly  are  these 
and  other  sulfur  reagents  used  by  chemists,  because  of  their 
cheapness  and  wide  range  of  adaptability,  that  they  may  be 
counted  among  the  most  important  tools  of  the  trade.     It  is 
therefore  from  the  point  of  view  of  possessing  ourselves  of  a 
working  knowledge  of  these  important  reagents  that  we  are 
next  to  undertake  the  further  study  of  sulfur  and  its  com- 
pounds. 

600.  The  Physical  Properties  of  Sulfur. — The  commonest 
form  of  free  sulfur  is  called  rhombic  sulfur  from  the  shape  of  its 
crystals.    It  is  light  yellow  in  color,  not  soluble  in  water,  slightly 
soluble  in  alcohol  or  ether,  but  very  soluble  in  some  of  its  own 
compounds,  such  as  carbon  disulfide,  CS2  (546).    Another  form 
of  sulfur,  monoclinic  sulfur,  which  is  also  named  from  the  shape 
of  its  crystals,  has  the  same  color  as  rhombic  sulfur  and  is 
soluble  in  the  same  solvents  but  to  a  different  degree.    At  95°, 
the  so-called  transition  point,  these  two  forms  of  sulfur  can  exist 
in  contact  with  each  other,  just  as  ice  and  water  can  at  o°.    If 
the  mixture  of  the  two  forms  is  heated  above  95°,  all  the  rhombic 
sulfur  becomes  monoclinic  sulfur.    On  the  other  hand,  if  the 
mixture  is  cooled  below  95°  all  the  monoclinic  sulfur  changes  to 

375 


376  Introduction  to  General  Chemistry 

rhombic  sulfur.  If  the  monoclinic  sulfur  is  cooled  quickly  to 
room  temperature  the  change  to  rhombic  sulfur  goes  on  very 
slowly,  just  as  the  change  of  white  to  red  phosphorus  goes  on 
slowly  at  ordinary  temperatures,  though  it  is  rapid  at  higher 
temperatures.  It  is  not  possible  to  heat  a  solid  substance  above 
its  melting-point  without  melting  it,  but  it  is  possible  to  heat  a 
solid  above  its  transition  point  to  another  solid  form  and  to 
determine  the  melting-point  of  the  first  form,  provided  the  opera- 
tion is  carried  on  very  rapidly.  In  this  way  it  has  been  found 
that  rhombic  sulfur  melts  at  114°.  Monoclinic  sulfur  melts 
at  119°. 

Still  another  form  of  sulfur,  plastic  sulfur,  is  prepared  by 
heating  sulfur  to  its  boiling-point  and  then  pouring  it  into  cold 
water.  Sulfur  so  treated  is  a  plastic  solid,  dark  amber  in  color, 
and  when  free  from  rhombic  sulfur,  which  may  be  formed  with 
it,  it  is  insoluble  even  in  carbon  disulfide.  In  time  it  changes  to 
rhombic  sulfur. 

With  rising  temperature  molten  sulfur  is  at  first  a  mobile, 
straw-colored  liquid,  but  at  about  160°,  curiously  enough,  it 
becomes  dark  brown  and  extraordinarily  viscous.  But  when  its 
temperature  has  reached  260°,  its  color  becomes  a  yellow  red, 
and  its  viscosity  again  decreases.  These  changes  are  due  to  the 
formation  of  allotropic  liquid  forms.  This  is  the  first  case  of  the 
kind  which  we  have  encountered. 

60 1.  Chemical  Properties  of  Sulfur. — Sulfur  unites  chemi- 
cally with  most  metals  and  non-metals.  Thus  it  not  only  forms 
sulfides  of  metals,  but  it  also  forms  oxides,  chlorides,  bromides, 
phosphides,  etc.  For  example,  when  it  is  rubbed  together  with 
mercury,  black  mercuric  sulfide  forms: 

Hg-hS->HgS. 

If  sulfur  is  heated  with  iron,  the  reaction  to  form  ferrous 
sulfide  begins  and  soon  gives  out  enough  heat  to  raise  the  whole 
mass  to  incandescence  (339).  The  oxides  of  sulfur  are  of  such 
importance  that  their  chemistry  will  be  described  later  in  detail. 
Chlorine  passed  over  melted  sulfur  forms  sulfur  monochloride, 


Sulfur  and  Its  Compounds  377 

S2C12.  The  latter  substance,  which  is  a  liquid  at  ordinary 
temperatures,  is  an  important  solvent  for  sulfur.  It  has  been 
produced  in  enormous  quantities  for  the  preparation  of  the 
terrible  mustard  gas  (695). 

602.  The  Sources  of  Sulfur. — Sulfur  is  found  in  large  deposits, 
supposedly  of  volcanic  origin,  in  Sicily,  Japan,  and  elsewhere. 
Of  these  deposits,  those  of  Sicily  supplied  most  of  the  world's 
trade  up  to  1904.    At  that  time  two  mines  in  Louisiana  and  one 
in  Texas  became  of  tremendous  importance,  owing  to  the  develop- 
ment of  the  Frasch  process  of  getting  out  the  sulfur.    The 
deposits  now  being  mined  there  are  several  hundred  feet  below 
the  surface.     Iron  pipes  are  driven  down  into  the  beds,  and 
superheated  steam  is  forced  down  to  melt  the  sulfur,  which 
collects  in  a  pool  at  the  end  of  the  pipe.    From  this  pool  it  is 
forced  to  the  surface  by  compressed  air.     Single  wells  deliver 
per  day  as  much  as  four  hundred  to  five  hundred  tons  of  sulfur, 
99 . 5  per  cent  pure,  and  keep  it  up  for  months  at  a  time.    The 
sulfur  issuing  at  the  surface  is  piped  into  bins,  where  it  is  spread 
out  in  thin  layers,  which  immediately  solidify.     In  this  way  great 
blocks  of  sulfur  are  built,  sometimes  as  much  as  150  feet  wide, 
250  feet  long,  and  60  feet  high.    The  greater  part  of  the  sulfur 
carried  to  the  coast  from  the  mines  is  transported  in  dump- 
bottom  cars,  much  as  coal  is  transported. 

A  very  considerable  quantity  of  sulfur  is  produced   as  a 
by-product  of  other  industries. 

603.  Rock  or  Roll  Sulfur,  Flowers  of  Sulfur. — Sulfur  from 
the  mines  or  industry  is  called  roll  or  rock  sulfur.     Sometimes 
the  vapor  of  sulfur  is  quickly  cooled  to  produce  a  powder  called 
flowers  of  sulfur.    It  is  usually  contaminated  with  traces  of 
sulfur ous  and  sulfuric  acid. 

604.  Commercial  Importance  of  Sulfur.— One  of  the  very 
strikingly  important  uses  of  sulfur  is  found  in  the  process  of 
vulcanizing  rubber.    Experts  tell  us  that  without  sulfur  there 
would  be  no  rubber  industry  worthy  of  the  name.    It  seems  that 
crude  rubber  is  stiff  and  hard  at  the  temperature  of  ordinary 
winter  weather,  and  at  moderate   summer  heat  it  becomes 
very  sticky.     But  after  vulcanizing,  a  treatment  with  sulfur, 


Introduction  to  General  Chemistry 

commonly  as  vapor  or  with  antimony  sulfide  (red  rubber),  it 
gains  greatly  in  strength  and  elasticity  and  becomes  remarkably 
indifferent  to  heat. 

Free  sulfur  is  an  important  constituent  of  many  germicide 
and  of  many  fungicide  mixtures.  Finely  divided  sulfur  is 
dusted  on  vines  and  hops.  Before  the  war  Europe  used  more 
than  one  hundred  thousand  tons  per  annum  for  this  purpose 
alone.  Some  sulfur  is  used  in  making  gunpowder,  the  so-called 
"  black  powder."  Enormous  quantities  of  sulfur  are  used  to 
make  sulfur  compounds  such  as  sulfur  dioxide,  sulfites,  sulfuric 
acid,  carbon  disulfide,  and  a  host  of  other  important  sub- 
stances. Before  the  war  the  annual  consumption  of  sulfur  in 
the  United  States  was  about  three  hundred  thousand  tons,  but 
during  the  war  it  was  reported  to  be  as  high  as  nine  hundred 
thousand  tons. 

605.  Hydrogen  Sulfide. — The  important  reagent  hydrogen 
sulfide  will  have  been  used  in  the  laboratory  by  the  time  the 
work  has  been  carried  to  this  point.    A  brief  description  of  its 
physical  properties  and  chemistry  has  already  been  given  (339) . 
Usually  generators  of  this  gas  are  available  in  every  laboratory. 
Many  special  types  have  been  designed,  but  in  principle  they 
are  much  like  the  hydrogen  generators  already  described  in 
294.    Ferrous  sulfide  and  hydrochloric  acid  are  almost  uni- 
versally used  to  charge  them. 

Hydrogen  sulfide  burns  in  air  to  form  water  and  sulfur 
dioxide.  If,  however,  the  flame  is  allowed  to  play  on  a  cool 
surface  it  deposits  sulfur.  Either  as  a  gas  or  dissolved  in  water 
it  is  a  very  powerful  poison,  small  doses  causing  insensibility  or 
even  death.  Fortunately  its  very  bad  odor  warns  of  its  presence, 
so  that  it  can  usually  be  avoided. 

606.  Water  Solution  of  Hydrogen  Sulfide.     Hydrosulfurous 
Acid. — One  volume  of  water  at  o°  dissolves  eighteen  volumes  of 
the  gas,  but  only  about  three  volumes  at  room  temperature.    At 
the  latter  temperature  and  atmospheric  pressure  the  saturated 
solution  is  approximately  one- tenth  molar  in  concentration. 
All  the  hydrogen  sulfide  may  be  removed  from  a  solution  by 
boiling  it. 


Sulfur  and  Its  Compounds  379 

The  solution  will  turn  neutral  litmus  faintly  pink.  Hydrogen 
sulfide,  or  hydrosulf urous  acid,  ionizes  in  stages  as  follows : 

H2S->H++HS~, 
HS--»H++S— . 

The  first  or  primary  dissociation  is  weaker  than  that  of  carbonic 
acid  or  of  dihydrogen  phosphate  ion,  H2P04~  (592).  The 
secondary  dissociation  is  extremely  small.  H2S  reacts  with 
sodium  hydroxide  to  form  two  products,  sodium  hydrogen  sulfide 
and  sodium  sulfide: 

H2S+NaOH->NaHS+H2O, 
NaHS+NaOH->Na2S+H2O. 

607.  Sodium  and  Ammonium  Sulfides,  Yellow  Ammonium 
Sulfide. — The  first  of  these  reactions  is  nearly  complete,  but  the 
second  is  far  from  being  so.  Crystalline  sodium  sulfide,  Na2S, 
can,  however,  be  prepared.  A  water  solution  of  sodium  hydrogen 
sulfide,  NaHS,  is  slightly  alkaline,  but  a  solution  of  the  other  salt 
is  strongly  alkaline.  These  are  the  results  which  we  would 
expect  from  the  weakness  of  the  acids  H2S  and  HS~  (436). 
Solutions  of  sodium  sulfide  find  important  use  as  a  depilatory 
for  the  removal  of  hair  from  the  skins  of  slaughtered  animals  and 
for  other  similar  purposes. 

The  ammonium  salts,  NH4HS  and  (NH4)2S,  are  prepared  by 
passing  hydrogen  sulfide  into  ammonium  hydroxide  solution. 
The  resulting  solution,  which  is  of  course  alkaline  in  reaction, 
is  extensively  used  in  chemical  analysis. 

Although  pure  ammonium  sulfide  solutions  are  colorless,  the 
ordinary  reagent  is  usually  yellow,  owing  to  the  presence  of 
dissolved  sulfur  formed  because  of  exposure  to  the  oxidizing 
influence  of  the  air.  Sulfur  dissolves  in  ammonium  sulfide 
forming  a  series  of  complex  ions  which,  because  of  their  similarity 
to  peroxide  ion,  are  called  persulfide  ions.  Ions  of  the  formula 
S2~~,  S3~~,  etc.,  are  known.  The  laboratory  reagent,  yellow 
ammonium  sulfide,  which  is  prepared  by  dissolving  sulfur  in 
ammonium  sulfide,  probably  contains  a  mixture  of  persulfides. 
Its  formula  is  usually  given  (NH4)2Sa;,  the  subscript  x  denoting 


380  Introduction  to  General  Chemistry 

an  undetermined  proportion  of  sulfur.     Upon  acidification  of 
this  reagent,  sulfur  is  precipitated  and  hydrogen  sulfide  given  off  : 

(NH4)2S*+  2HCl->  2NH4C1+H2S*, 


608.  The  Precipitation  of  Sulfides.  —  The  precipitation  of 
sulfides  from  solution  is  a  very  important  procedure  in  making 
an  analysis  of  a  substance  of  unknown  composition.    The  pro- 
cess of  a  general  metal  analysis  depends  on  the  separation  of 
metals  from  the  mixture  in  order  that  they  may  be  recognized 
by  their  characteristic  reactions.    Thus  if  we  have  in  solution 
copper  sulfate  and  zinc  sulfate,  we  may  separate  the  metals  as 
follows:  First,  the  solution  is  made  acid  (o.  25  to  o.  5  N)  with 
hydrochloric  acid,  and  hydrogen  sulfide  is  passed  in.     Copper 
sulfide  alone  is  precipitated,  because  the  excess  of  acid  makes 
the  concentration  of  sulfur  ion  so  small  (432)  that  only  the 
extremely  insoluble  copper  sulfide  forms  in  sufficient  amount  to 
exceed  its  molecular  solubility  (445)  .    Next,  the  filtrate,  which  is 
practically  free  from  copper,  is  made  alkaline  with  ammonium 
hydroxide.     Immediately  the  separation  of  the  white  zinc  sulfide, 
from  solution  begins  and  is  carried  to  completion  by  the  further 
addition   of   ammonium   sulfide.     In  practice   the   extremely 
insoluble  sulfides  of  mercury,  lead,  bismuth,  cadmium,  arsenic, 
antimony,  tin,  etc.,  if  present,  are   brought  down  with  the 
copper  sulfide,  and  then  further  separation  of  these  metals  from 
each  other  follows.    With  zinc  sulfide  appear  the  sulfides  of  iron, 
cobalt,  nickel,  and  manganese,  together  with  certain  insoluble 
hydroxides.    The  separation  by  means  of  hydrogen  sulfide  is 
therefore  one  of  the  fundamental  processes  in  the  analysis  of 
metals. 

609.  Aqueous  Hydrogen  Sulfide  as  a  Reducing  Agent.— 
Aqueous  hydrogen  sulfide  is  a  good  reducing  agent  by  virtue  of 
the  easily  oxidizable  sulfur  ion,  S      ,  which  it  furnishes  (503). 
Thus  when  it  is  oxidized  by  the  air  (339)  the  fundamental 
change  may  be  represented  as  follows: 

28—  -+O2->2S+2O  —  . 


Sulfur  and  Its  Compounds  381 

The  hydrogen  ions  which  had  been  associated  with  the  sulfur  ion 
before  the  change  unite  with  the  oxygen  ion  produced,  forming 
water. 

Since  hydrogen  sulfide  is  acted  upon  in  water  solution  by 
such  mild  oxidizing  agents  as  the  oxygen  of  the  air,  it  is  not 
surprising  to  find  that  it  is  attacked  by  very  powerful  agents 
such  as  acid  permanganate  (343)  or  acid  dichromate  solutions 
(346).  If  dilute  sulfuric  acid  is  added  to  potassium  permanga- 
nate solution,  and  hydrogen  sulfide  is  passed  into  the  mixture, 
the  intense  color  of  the  permanganate  soon  disappears,  and  in 
its  place  we  have  the  nearly  colorless  solution  of  .manganous 
sulfate.  The  sulfur  ion  may  be  oxidized  to  free  sulfur  or  to 
sulfite  ion  or  sulfate  ion  according  to  the  experimental  conditions. 
If  free  sulfur  is  the  main  product,  as  shown  by  a  fine,  almost 
white  suspension  in  solution,  the  fundamental  change  (561)  in 
charges  of  the  atoms  undergoing  oxidation  and  reduction  is  as 
follows  : 


and  the  final  equation  is 

2KMnO4+5H2S+3H2SO4->2MnS04+K2S04+5S+8H2O. 

When  hydrogen  sulfide  is  led  into  potassium  dichromate  solu- 
tion mixed  with  a  strong  acid,  sulfuric  acid  for  instance,  the 
characteristic  color  change  from  orange  to  green  or  violet  of  the 
latter  reagent  on  reduction  is  soon  observed.  In  this  case  also 
free  sulfur,  sulfurous  acid,  or  sulfuric  acid  may  be  produced  by 
the  reaction.  If  the  fine,  almost  white,  precipitate  of  sulfur  is 
observed  the  fundamental  equation  is 


->  2Cr++++3S, 

and  the  final  equation  is 

K2Cr207+4H2S04+3H2S->K2S04+Cr2(S04)3+3S+7H20. 

610.  Sulfur  Dioxide.  —  Sulfur  dioxide  is  a  sharp-odored, 
colorless  gas  formed  when  sulfur  or  sulfides  are  burned  in  air. 
It  is  often  found  in  volcanic  gases  and  is  to  be  noticed  in  the  air 


382  Introduction  to  General  Chemistry 

about  towns,  since  much  of  the  coal  burned  contains  traces  of 
sulfur.  A  99  per  cent  pure  gas  is  made  on  a  huge  scale  by  burning 
sulfur,  but  an  enormous  supply  is  also  obtained  by  the  burning 
of  pyrite,  FeS2,  one  of  the  richest  of  the  sulfur  ores: 

4FeS2+ 1  iO2  ->  2Fe2O3+8SO2. 

Since  the  gas  is  very  easily  liquefied  it  is  generally  marketed  in 
that  form  and  stored  in  iron  tanks,  the  material  of  which  it 
does  not  attack.  At  ordinary  temperatures  the  pressure  in 
these  tanks  is  about  2 . 5  atmospheres.  The  gas  dissolves  easily 
water,  about  70  volumes  to  one  of  water  at  o°  and  about  40 
volumes  at  room  temperature.  A  saturated  solution  at  20° 
and  one  atmosphere  pressure  is  about  a  i .  6  molar  solution. 
The  sulfurous  acid  of  commerce  contains  about  6  per  cent  of 
sulfur  dioxide. 

611.  Water  Solution  of  Sulfur  Dioxide,  Sulfurous  Acid. — In 
water  solution  sulfur  dioxide  forms  sulfurous  acid  (340)  just  as 
carbon  dioxide  forms  carbonic  acid  (152).    Sulfurous  acid  is  a 
moderately  weak  dibasic  acid.    Its  reducing  power  has  already 
been  described  (340).    On  this  account  it  is  used  on  an  enormous 
scale  commercially  as  "antichlor,"  that  is,  to  remove  the  residual 
hypochlorous  acid  left  after  the  bleaching  action  of  chlorine. 
Sulfurous  acid  is  also  used  to  bleach  substances  which  chlorine 
rots,  such  as  wool,  silk,  etc.     Sulfurous  acid  and  sulfites  are  used 
in  breaking  up  the  fibers  of  wood  pulp  preliminary  to  paper- 
making.     Sulfur  dioxide  is  also  used  in  the  preservation  of  food, 
particularly  dried  fruits. 

612.  Salts  of  Sulfurous  Acid,  the  Sulfites. — Sulfurous  acid 
forms  both  acid  and  normal  salts.     Most  of  the  acid  salts  are 
soluble,  but  of  the  normal  salts  only  the  sodium,  potassium,  and 
ammonium  sulfites  are  soluble  in  water.    Like  carbonates  (461) 
sulfites  interact  with  strong  acids  to  form  the  correspondingly 
strong  acid  salt  and  the  weak  and  unstable  acid,  which  decom- 
poses, giving  off  the  acid  anhydride: 

BaS03+  2HC1  ->  BaCl2+H2SO3, 
H2SO3->H2O+SO3. 


Sulfur  and  Its  Compounds  383 

This,  of  course,  means  that  sulfites,  which  are  only  moderately 
insoluble,  are  dissolved  extensively  by  acids,  particularly  if 
the  sulfur  dioxide  is  removed  from  solution  by  boiling  (463). 

613.  Sulfur  Trioxide.    Fuming  Sulfuric  Acid. — Sulfur  tri- 
oxide,  S03,  the  anhydride  of  sulfuric  acid,  is  a  choking  gas.     It 
dissociates  into  sulfur  dioxide  and  oxygen  more  and  more  with 
increasing  temperature;  at  400°  this  dissociation  is  only  about 
2  per  cent  of  the  whole;  at  500°  it  is  about  9  per  cent;  at  1000° 
it  is  virtually  complete.    Sulfur  trioxide  can  be  condensed  into 
needle-shaped  crystals  which  melt  at  18.8°  and  form  a  liquid 
which  may  be  distilled  undecomposed  at  44 . 8°.    In  the  presence 
of  a  trace  of  moisture  which,  of  course,  forms  sulfuric  acid,  sulfur 
trioxide  changes  to  a  white,  asbestos-like  mass  which  can  be 
melted  at  50°.    This  substance  probably  consists  of  molecules, 
each  one  of  which  contains  several  sulfur  trioxide  molecules. 
Substances  the  molecules  of  which  are  made  up  of  small  mole- 
cules of  a  simpler  substance  are  called  polymers.     Sulfur  trioxide 
is  very  hygroscopic.     When  thrown  into  water  it  reacts  violently 
to  form  sulfuric  acid. 

Sulfur  trioxide  dissolves  in  sulfuric  acid  to  form  a  product 
called  fuming  sulfuric  acid  or  oleum. 

614.  Sulfuric  Acid. — Sulfuric  acid  has  a  clear  title  to  being 
the  most  important  chemical  used  in  industry.     Very  few 
chemical  industries  exist  which  do  not  employ  this  acid  in  some 
process.    During  normal  times  the  United  States  produced 
about  three  and  a  half  million  tons  annually,  but  during  the 
Great  War  the  yearly  output  is  reported  to  have  exceeded 
six  million  tons  (see  618).    This  statement  does  not  include 
enormous  amounts  made  and  used  without  being  put  on  the 
market. 

615.  The  Manufacture  of  Sulfuric  Acid. — The  problems  of 
the  manufacture  of  sulfuric  acid  are  much  like  those  of  the 
synthesis  of  ammonia  (520).    At  low  temperatures  the  union  of 
sulfur  dioxide  with  oxygen  to  form  sulfur  trioxide,  the  anhydride 
of  sulfuric  acid,  will  go  to  completion,  but  too  slowly  for  the 
process  to  be  profitable.    If  higher  temperatures  are  used  the 
rate  of  the  reaction  is  increased,  but  equilibrium  is  established 


384  Introduction  to  General  Chemistry 

with  less  of  sulfur  trioxide  formed  the  higher  the  temperature 
used.  The  manufacture  of  the  acid  has  been  made  successful 
by  the  use  of  catalysts,  which  hasten  the  reaction  at  low  tempera- 
tures. Two  general  methods  are  used.  The  one  called  the 
chamber  process  relies  on  a  gas  catalyst  and  allows  the  reaction 
to  come  to  completion  in  huge  reaction  chambers  of  lead.  The 
other,  called  the  contact  process,  employs  a  solid  catalyst,  on 
the  sur&ce  of  which  the  reaction  proceeds. 

616.  The  Chamber  Process. — In  the  first  method,  which  is 
the  older,  the  catalyst  is  a  mixture  of  nitric  oxide  and  nitrogen 
tetroxide.  Sulfur  dioxide,  supplied  usually  from  the  burning 
of  sulfur  or  pyrite,  is  mixed  with  the  right  amount  of  air  and  is 
then  swept  over  pots  which  contain  niter  and  sulfuric  acid. 
After  receiving  the  nitric  acid  fumes  from  the  latter  the  mixed 
gases  are  sent  to  the  so-called  Glover  tower  to  receive  more  of 
the  catalyst,  which  has  been  recovered  from  previous  charges  in 
a  manner  which  will  be  described.  The  reaction  begins  in  the 
Glover  tower,  but  from  this  tower  the  gases  are  forced  into  the 
great  lead  chambers  in  which  the  reaction  is  completed.  Jets 
of  cold  water  or  dilute  sulfuric  acid  in  the  chambers  dissolve  the 
sulfur  trioxide,  forming  a  moderately  concentrated  sulfuric  acid, 
called  chamber  acid: 

S03+H2O->H2SO4. 

The  actual  mechanism  by  which  the  catalyst  acts  is  not  com- 
pletely understood.  The  mixture  of  NO  and  NO2  undoubtedly 
contains  a  little  N203  (554)  in  equilibrium  with  the  former.  If 
nitrous  acid  anhydride  is  regarded  as  the  active  agent,  the  reac- 
tion may  be  represented  by  the  following  equation: 

N2O3+ SO2  ^  SO3+  2NO. 

The  nitric  oxide  left  combines  with  more  oxygen,  and  the  regen- 
erated oxides  are  again  able  to  oxidize  more  sulfur  dioxide. 

The  gases  left  in  the  chambers  are  next  forced  through  the 
Gay  Lussac  tower.  The  latter  is  filled  with  broken  brick,  over 
which  concentrated  sulfuric  acid  constantly  trickles.  In  this 
both  sulfur  trioxide  and  nitric  oxide  dissolve,  the  first  to  form 


Sulfur  and  Its  Compounds  385 

fuming  sulfuric  acid  (613)  and  the  second  to  form  nitrosyl 
sulfuric  acid: 

4NO+O2-f  4H2SO4  ->  4NO  •  HS04+ 2H2O. 

The  mixture  is  drawn  off  and  carried  to  the  top  of  the  Glover 
tower.  From  this  point  it  is  allowed  to  run  down  over  the  brick 
with  which  the  tower  is  filled.  As  it  does  so  it  meets  a  stream 
of  dilute  acid.  The  dilution  of  the  nitrosyl  sulfuric  acid  effects 
its  decomposition,  and  oxides  of  nitrogen  are  given  off  to  the 
gases  entering  the  tower  from  the  niter  pots: 

2NO-HSO4+H2O->  2H2SO4+ NO2+ NO. 

The  nitric  oxide  is  of  course  changed  into  higher  oxides  by  the 
oxygen  of  the  air,  so  that  the  catalyst  is  being  continually 
re-formed.  The  mixed  gases  then  go  on  the  next  trip  through 
the  lead  chambers.  A  little  fresh  catalyst  is  furnished  con- 
tinually from  the  niter  pots  to  make  up  a  small  unavoidable  loss 
in  the  process.  The  object  of  the  broken  brick  in  the  two  towers 
is  to  afford  a  large  surface  of  contact  for  the  gases.  The  yield 
of  sulfuric  acid  is  about  95  per  cent  of  the  amount  theoretically 
possible.  The  chamber  acid  is  of  a  concentration  of  about 
78  per  cent  H2S04  and  22  per  cent  water.  More  concentrated 
acid  is  made  by  subsequent  evaporation. 

The  following  diagram,  a  so-called  flow  sheet,  shows  by 
means  of  arrows  how  the  various  materials  entering  into  and 
formed  in  the  chamber  process  are  related  to  one  another: 

Sulfur  Air 

So2    NA 

H2O     SO3        NO        Air 
Dil.      H2SO4      NA        Cone.  H2SO4 
Cone.  H2SO4     Dil.  H2SO4  NO-HSQ4 

v  v 

H2S04  NA 


386  Introduction  to  General  Chemistry 

Sulfur,  air,  water,  and  the  oxides  of  nitrogen  are  the  start- 
ing materials.  Of  these  the  oxides  of  nitrogen,  which  act 
catalytically,  are  largely  recovered  in  the  process  and  used  over 
and  over  again.  The  small  deficiency  caused  by  the  failure  to 
recover  them  completely  is  made  up  by  gases  from  concentrated 
nitric  acid.  i 

617.  The  Contact  Process. — The  more  recent  method  of 
making  sulfuric  acid  is  the  so-called  contact  process.  Sulfur 
dioxide  is  made  to  unite  with  oxygen  on  the  surface  of  a  catalyst 
which  can  be  iron  oxide,  nickel  oxide,  brick,  quartz,  etc. 
Platinum  is  an  especially  good  agent  for  this  purpose,  since  it 
allows  the  reaction  to  be  carried  on  at  temperatures  as  low  as 
450-500°.  Other  catalysts  are  reported  to  be  in  use.  If  platinum 
is  employed  the  gases  must  be  carefully  freed  from  arsenic,  since 

TABLE  XXV 


Be" 
(Degrees  Baume1) 

Tw 
(Degrees  Twadell) 

Specific  Gravity 

Percentage  of 
Sulfuric  Acid 

i°° 

60 
66 

105° 
141 
167 

1-53 
1.71 
1.84 

62 
78 
93 

this  substance,  which  is  commonly  present  in  pyrite,  renders 
the  platinum  ineffective.  The  sulfur  trioxide  formed  is  dis- 
solved in  concentrated  sulfuric  acid,  and  water  is  added  to  the 
product  to  make  the  more  dilute  acids.  The  contact  method  has 
the  advantage  that  much  more  concentrated  solutions  of  sulfuric 
acid  may  be  prepared  directly  by  it  than  by  the  chamber  process; 
but  in  spite  of  this  advantage  and  the  greater  simplicity  of  the 
method,  the  chamber  process  is  still  the  more  commonly  used, 
owing  to  many  skilful  improvements  in  the  details  of  manu- 
facture. 

618.  The  Baume  and  Twadell  Scales  of  Specific  Gravity. — 
In  the  market  reports,  etc.,  the  concentration  of  sulfuric  acid  is 
usually  given  in  terms  of  either  of  two  arbitrary  scales  of  specific 
gravity:  in  England  and  the  United  States  the  Baume  (abbre- 
viation Be)  and  on  the  Continent  the  Twadell  (abbreviation  Tw). 
Table  XXV  gives  some  data  for  the  temperature  of  15°. 


Sulfur  and  Its  Compounds  387 

The  "  concentrated  sulfuric  acid"  of  the  laboratory  is  usually 
understood  to  be  66°  Be.  The  annual  production  of  sulfuric  acid 
is  given  in  terms  of  66°  acid  (614).  The  so-called  chamber  acid 
(6  1  6)  is  usually  60°  Be.  The  specific  gravity  of  other  liquids 
heavier  than  water  is  also  given  on  these  scales.  Another 
Baume  scale  is  used  for  liquids  lighter  than  water.  Complete 
tables  showing  the  relationship  between  the  specific-gravity 
scales  will  be  found  in  most  collections  of  chemical  tables. 

619.  Physical  Properties  of  Sulfuric  Acid.  —  Pure  sulfuric  acid 
is  a  heavy,  oily  liquid  often  called  oil  of  vitriol.    Like  hydro- 
chloric acid  it  forms  with  water  a  mixture  of  constant  boiUng- 
point  (251).    At  one  atmosphere  pressure  the  compositroti  of 
this  mixture  is  98.4-6  per  cent  acid.     It  boils  at  338°.    Pure 
sulfuric  acid  boils  at  290°.    The  vapor  of  sulfuric  acid  consists 
almost  entirely  of  sulfur  trioxide  and  water. 

Sulfuric  acid  is  very  hygroscopic  and  therefore  a  good  drying 
agent.  Since  the  98  .  4  per  cent  acid  does  not  give  off  fumes  at 
ordinary  temperatures,  a  little  is  often  kept  in  storage  vessels, 
such  as  desiccators,  to  keep  the  air  dry.  If  a  moist  substance  is 
put  in  such  a  vessel  it  will  continuously  lose  water  to  the  air, 
and  as  the  water  vapor  is  constantly  being  taken  up  by  the 
sulfuric  acid  the  substance  is  soon  dry.  The  great  heat  which  is 
given  out  as  sulfuric  acid  is  diluted  with  water  has  already  been 
mentioned.  Students  have  been  warned  (93)  that  there  is 
danger  from  spattering  whenever  sulfuric  acid  is  mixed  with 
water.  On  this  account  the  acid  is  always  poured  very  slowly 
into  water,  with  gentle  stirring  in  order  that  the  heavy  layer  may 
be  well  mixed  into  the  lighter  layer.  The  process  of  pouring 
water  into  sulfuric  acid  is  very  likely  to  result  in  a  violent 
spattering  of  the  hot  and  corrosive  acid. 

620.  Chemical  Properties  of  Sulfuric  Acid.  —  Dilute  sulfuric 
acid  is  a  dibasic  acid  a  little  weaker  than  the  strongest  acids: 


Even  its  secondary  dissociation  is  that  of  a  moderately  strong 
acid.    Like  other  acids,  dilute  sulfuric  acid  reacts  with  metals 


388  Introduction  to  General  Chemistry 

above  hydrogen  in  the  electromotive  series  (492)  to  give  off 
hydrogen,  but  the  sulfate  ion  is  usually  very  slowly  affected  by 
reducing  agents. 

On  account  of  its  high  boiling-point  concentrated  sulfuric 
acid,  contrary  to  the  usual  rule,  can  be  made  to  displace  stronger 
acids  from  their  salts,  since  the  acid  formed  by  the  partial  double 
decomposition  of  sulfuric  acid  with  the  salt  is  distilled  off  at 
temperatures  at  which  the  sulfuric  acid  is  but  little  volatile. 
In  this  way  even  the  insoluble  silver  chloride  may  be  dissolved, 
though  the  action  is  very  slow: 

AgCl(solid)^AgCl(dissolved)^Ag++Cl- 


It  It 

AgHS04    HC1 

Concentrated  sulfuric  acid,  unlike  the  dilute,  is  an  oxidizing 
agent  as  well  as  an  acid,  though  its  power  to  oxidize  is  not  as 
strong  as  is  that  of  nitric  acid.  Some  of  the  important  and 
characteristic  reactions  will  be  discussed  shortly  (622). 

If  concentrated  sulfuric  acid  is  put  on  filter  paper  the  latter 
will  be  found  to  blacken.  This  is  a  typical  behavior  of  sulfuric 
acid  with  any  carbon  compound  from  which  it  can  take  water 
from  combination.  Thus  sugar,  CiaH^Ou,  is  charred  by  sulfuric 
acid.  For  the  same  reason  sulfuric  acid  burns  the  skin.  These 
reactions  are  in  reality  more  complex  than  simple  dehydrations. 

621.  Concentrated  Sulfuric  Acid  and  Metals.  —  Sulfuric  acid 
does  not  attack  silver  or  gold,  but  the  impure  acid  does  attack 
platinum.    In  the  cold  it  does  not  react  with  the  less  active 
metals  like  copper,  but  does  do  so  when  hot,  forming  the  sulfate 
of  the  metal  and  sulfur  dioxide.    The  more  active  metals  like 
zinc  are  attacked  by  the  warm  acid  with  the  liberation  of  S02; 
but  sulfur  and  hydrogen  sulfide  may  be  formed  according  to 
experimental  conditions.    That  sulfuric  acid  in  the  absence  of 
air  does  not  attack  iron  is  of  great  technical  importance,  since 
the  latter  material  is  widely  used  for  containers  of  this  acid. 

622.  Oxidation  by  Concentrated  Sulfuric  Acid.—  The  solu- 
tion of  metals  in  sulfuric  acid  is  of  course  the  result  of  the  oxidizing 


Sulfur  and  Its  Compounds  389 

action  of  the  acid.  Table  XXVI  states  briefly  the  charges 
which  the  sulfur  atom  carries  in  sulfuric  acid  and  its  reduction 
products,  as  well  as  in  the  other  common  compounds  which  we 
have  studied.  Apparently  a  sulfur  atom  in  sulfuric  acid  must 
gain  two  electrons  to  become  sulfite,  six  to  become  free  sulfur, 
and 'eight  to  become  hydrogen  sulfide. 

TABLE  XXVI 

Compounds  Sulfur  Atoms 

S03,  H2SO4,  and  sulfates S6+ 

SO2,  H2SO3,  and  sulfites S<+ 

Free  sulfur S° 

H2S S-" 

Of  the  possible  reduction  products  of  sulfuric  acid,  S  and 
H2S  are  capable  of  further  reaction  with  excess  acid.  Thus 
H2S  bubbled  through  concentrated  H2SO4  forms  sulfur  and  the 
unstable  sulfurous  acid: 

H2S-f-H2SO4->S+H2O+H2SO3. 

Sulfur  warmed  with  concentrated  H2SO4  reduces  the  acid,  form- 
ing S02: 

S+  2H2SO4  ->  SO2+  2H2SO3. 

As  in  the  case  of  nitric  acid  (557),  so  with  sulfuric  acid,  the 
concentration  of  the  acid,  the  temperature  of  the  reaction,  the 
fineness  of  division  of  the  reducing  agent  if  it  is  a  solid,  and 
whether  or  not  the  acid  is  present  in  excess,  all  affect  the  nature 
of  the  products. 

In  the  case  of  zinc  the  fundamental  reactions  by  which 
each  of  the  possible  products  is  formed  are  as  follows: 

for  SO2,  Zn+S6+  ->Zn+++S«+, 
for  S,  3Zn-f-S6+  ->3Zn+++S°, 
forH2S,  4Zn+S6+->4Zn+++S--, 

and  the  final  equations  are 

Zn+  2H2SO4  ->  ZnSO4+  S02-f  2H2O, 
3Zn+4H2SO4->3ZnSO4+S+4H20, 
4Zn+  sH2SO4  ->  4ZnSO4+H2S+4H2O. 


3QO  Introduction  to  General  Chemistry 

« 

623.  Action  of  Non-Metallic  Reducing  Agents  on  Concen- 
trated Sulfuric  Acid.  —  The  non-metallic  reducing  agents  show 
the  same  differences  in  behavior  with  concentrated  sulfuric  acid 
as  do  the  metals.    Thus  hydrogen  bromide  reduces  sulfuric  acid 
to  form  sulf  urous  acid  and  bromine,  while  hydriodic  acid  reduces 
sulfuric  acid  to  sulfur  and  hydrogen  sulfide,  iodine  and  water 
being  formed  at  the  same  time  (341).    The  action  of  sulfur  and 
hydrogen  sulfide  respectively  lias  just  been  discussed  (622). 
Charcoal  (carbon)  reduces  the  acid  to  sulfurous  acid  and  is  itself 
oxidized  to  carbon  dioxide: 

C+  2H2SO4  ->  CO2+  2H2SO3. 

624.  Sulfates.  —  Solutions  of  the  acid  sulfates  are  acid  to 
litmus,  owing  to  the  fact  that  the  HS04~  ions  are  very  largely 
dissociated  in  solutions  of  even  moderate  dilution.     Sulfates 
as  a  rule  are  easily  soluble  in  water,  but  those  of  barium,  stron- 
tium, calcium,  and  lead  are  little  soluble.    Under  dry  heat  the 
acid  sulfates  of  sodium  and  potassium  first  lose  water,  forming 
the  so-called  pyrosulfates.     These  on  further  heating  change 
into  the  normal  sulfates  with  the  loss  of  sulfur  trioxide: 

2Na(HSO4)2  ->  Na2S2O7+H3O, 
Na2S2O7  ->  Na2SO4+  SO3. 

Normal  sulfates  of  the  alkalies  are  very  resistant  to  heat.  The 
others  break  down,  forming  the  corresponding  oxides  and  sulfur 
trioxide: 

Fe2(S04)3-»Fe2O3+3S03. 

625.  Sodium    Thiosulfate.  —  One    other    sulfur    compound 
should  be  mentioned.    This  is  sodium  thiosulfate,  the  "hypo" 
used  in  photography  to  "fix"  negatives.     If  sulfur  is  boiled  in  a 
solution  of  sodium  sulfite  it  will  soon  be  seen  to  dissolve.    Upon 
evaporation  the  solution  gives  clear,  colorless  crystals  of  sodium 
thiosulfate: 


This  salt  is  extremely  soluble  in  water.    At  o°,  21  7  parts  dissolve 
in  100  parts  by  weight  of  water.    It  will  be  remembered  as  one 


Sulfur  and  Its  Compounds  391 

of  the  substances  used  to  show  the  phenomenon  of  supersatura- 
tion  of  a  solution  (123).  When  solutions  of  hypo  are  made,  acid 
sulfur  is  deposited  and  sulfurous  acid  forms: 

Na2S203+  2HC1  ->  2NaCl+H2SO3+ S. 

Sodium  thiosulfate  is  manufactured  in  immense  quantities  for 
purposes  similar  to  those  for  which  sodium  sulfite  is  used.  The 
reaction  with  chlorine  is  as  follows: 

Na2S2O3+4Cl2+ 5H2O  ->  2NaCl+  2H2SO4+  6HC1. 

Sodium  thiosulfate  reacts  differently  with  iodine,  forming  sodium 
tetrathionate  and  sodium  iodide: 

2Na2S203+I2  ->  2NaI+Na2S4O6. 

626.  The  Silver  Thiosulfate  Complex  Ion. — If  a  solution  of 
sodium  thiosulfate  is  added  to  a  solution  of  silver  nitrate  a  white 
precipitate  of  silver  thiosulfate  appears: 

2AgN03+Na2S203  ->  Ag2S203+  2NaNO3. 

The  precipitate  dissolves  very  easily  in  excess  of  the  hypo  solu- 
tion, forming  complex  silver  thiosulfate  ions,  Ag2S406~~,  and 
secondarily  some  of  the  sodium  silver  thiosulfate  molecules : 

Ag2S2O3-fNa2S2O3->Na2Ag2S4O6. 

If  we  add  sodium  chloride  to  the  resulting  solution  no  precipitate 
of  silver  chloride  will  appear.  Or  if  we  treat  silver  chloride  with 
sodium  thiosulfate  solution  the  former  will  dissolve.  So  little 
dissociated  is  the  complex  silver  thiosulfate  ion  that  even  silver 
bromide,  which,  is  less  soluble  than  silver  chloride,  is  dissolved 
by  the  hypo  solution. 

627.  The  Action  of  Hypo  as  a  Fixing  Solution. — The  coating 
of  a  photographic  plate  or  film  which  is  sensitive  to  light  is 
ordinarily  made  of  gelatine  and  silver  bromide.    After  the  plate 
has  been  "exposed"  and  "developed"  the  parts  acted  upon  by 
light  are  black  coatings  of  finely  divided  silver.    The  unchanged 
part  of  the  "negative"  is  still  coated  with  the  light-yellow  silver 
bromide,  which  must  be  removed  before  the  former  is  again 


39 2  Introduction  to  General  Chemistry 

exposed  to  light,  or  the  whole  plate  will  become  evenly  black 
and  the  picture  be  obliterated.  Accordingly  the  unchanged  silver 
bromide  is  dissolved  in  the  hypo  bath,  and  the  soluble  sodium 
silver  thiosulfate  thus  formed  is  washed  away. 

628.  Other  Sulfur  Compounds. — The  chemistry  of  sulfur  and 
its  compounds  has  not  been  exhausted  by  any  means  by  the  dis- 
cussions of  this  chapter.  The  treatment  here  has  been  limited 
to  the  more  important  properties  of  the  commonest  sulfur  com- 
pounds; but  there  are  numerous  others,  some  of  which  are  of 
technical  importance.  Of  the  latter  class  are  persulfuric  acid, 
H2S2Os  permonosulfuric  or  Caro's  acid,  H2SOS;  and  hyposulfur- 
ous  acid,  H2S204,  together  with  their  salts.  The  first  two  of 
these  acids  are  closely  related  to  hydrogen  peroxide.  Their 
graphic  formulae  (323),  together  with  that  of  sulfuric  acid,  are 
thought  to  be  as  follows: 


HO-SO2  HO-SO2  HO-S02 

I  I 

o  o 


kJV 

A 


0  O  H 

I  I 

HO-SO2.  H 

Persulfuric  acid  Permonosulfuric  acid  Sulfuric  acid 

They  and  their  salts  are  powerful  oxidizing  agents.  The  Zn 
salt  of  hyposulfurous  acid  is  readily  formed  by  dissolving  zinc 
dust  in  sulf urous  acid : 

Zn+  2H2SO3  ->  ZnS2O4+  2H2O. 

The  sodium  salt  Na2S2O4,  known  in  commerce  as  sodium 
hydrosulfite,  is  a  powerful  reducing  agent  and  has  strong  bleach- 
ing action  on  vegetable  colors.  It  is  used  for  bleaching  soap,  etc. 
Five  other  acids  of  theoretical  but  at  present  little  practical 
importance  are  the  following:  dithionic  acid,  H2S2O6;  trithionic 
acid,  H2S3O6;  tetrathionic  acid,  H2S4O6;  pentathionic  acid, 
hexathionic  acid,  H2S606. 


CHAPTER  XXV 
CARBON  AND  CARBON  COMPOUNDS:   ORGANIC  COMPOUNDS.    I 

629.  Review. — Some  relatively  simple  reactions  of  carbon 
and   carbon   compounds  have   already   been   studied.    When 
carbon  is  burned  in  oxygen,  the  gas  carbon  dioxide  is  formed 
and  may  be  identified  by  its  reaction  with  limewater  to  form  a 
white  precipitate  of  calcium  carbonate  (18,  151).    The  burning 
of  other  carbon  compounds  in  excess  of  air  is  found  to  produce 
the  same  gas  (20).     Charcoal  heated  with  copper  oxide  reduces 
the  latter,  forming  free  copper  and  carbon  dioxide  (328).    The 
commercial  production  of  phosphorus  depends  on  the  reducing 
action  of  carbon  on  calcium  phosphate  mixed  with  sand  (580). 
Carbon  monoxide  is  also  a  good  reducing  agent  at  high  tempera- 
tures (329).    When  carbon  dioxide  dissolves  in  w;ater  the  weak 
and  unstable  carbonic  acid  is  formed  (152,  285,  449).    The 
bicarbonate  and  carbonate  of  sodium  and  of  potassium  are 
common   reagents   (161,    162,  448).    Limestone   and   marble, 
different  forms  of  calcium  carbonate,  are  used  on  a  huge  scale 
for  the  preparation  of  both  lime  (150)  and  carbon  dioxide  (163). 
The  precipitation  of  carbonates  is  governed  by  the  usual  condi- 
tions which  control  the  precipitation  of  little  soluble  salts  of 
weak  acids  (448).     All  carbonates  are  soluble  in  strong  acids 
(461).    Another  common  reagent,  acetic  acid,  is  also  a  carbon 
compound  (157).     It  is  a  colorless  liquid,  of  sour  odor,  miscible 
with  water  in  all  proportions.     It  is  used  mainly  as  a  weak  and 
stable  acid  (424). 

630.  Physical  Properties  of  Carbon. — The  free  element  car- 
bon appears  in  several  allotropic  forms.    An  impure  form  of 
great  industrial  importance  is  the  non-volatile  portion  of  coal 
left  after  the  latter  has  been  subjected  to  processes  which  free 
it  of  volatile  matter.    This  material  is  called  coke.     It  is 
employed  as  a  fuel  directly  but  is  also  used  for  making  different 
forms  of  fuel  gas,  which  are  usually  mixtures  of  carbon  monoxide 

393 


394  Introduction  to  General  Chemistry 

with  other  gases.  Coke  is  also  used  for  the  reduction  of  oxide 
ores  such  as  the  important  iron  ore  hematite,  Fe2O3.  Thus 
coke  is  a  necessary  material  for  many  industries. 

Graphite  is  another  form  of  carbon.  It  occurs  to  some  extent 
in  nature  but  is  now  produced  in  enormous  quantities  by  heating 
coke  in  the  electric  furnace  (Acheson  process).  It  is  formed 
in  slippery  black  scales  which  serve  as  an  excellent  lubricant 
for  parts  exposed  to  high  temperatures  and  for  wood  surfaces 
which  must  rub  together.  Since  graphite  conducts  electricity 
it  is  an  important  material  for  electrodes,  particularly  for  the 
preparation  of  substances  like  chlorine.  Mixed  with  clay, 
graphite  is  the  "lead"  of  lead  pencils. 

The  diamond  is  pure  crystalline  carbon.  The  natural 
crystals  are  cut  into  many-faced  shapes  in  order  to  bring  out 
the  power  of  this  material  to  reflect  light.  Diamonds  were 
first  produced  artificially  by  Moissan.  He  took  advantage  of 
the  fact  that  carbon,  though  insoluble  in  all  ordinary  solvents, 
is  soluble  in  molten  iron,  and  dissolved  graphite  in  the  latter 
substance.  He  then  chilled  the  mass  quickly  by  plunging  it 
into  molten  lead.  The  outside  layer  of  iron  was  first  to  cool, 
and  as  it  shrank  in  so  doing  it  put  the  inner  part  under  great 
pressure.  When  the  entire  mass  was  cooled  tiny  diamonds, 
about  0.5  mm.  in  diameter  at  the  most,  were  found  in  the 
interior.  The  diamond  has  the  distinction  of  being  the  hardest 
substance  known  and  is  consequently  an  important  abrasive 
and  cutting  material.  Black  diamonds,  which  have  no  orna- 
mental value,  are  set  in  the  cutting  surfaces  of  rock  drills,  while 
small  diamond  chips  are  used  to  cut  glass. 

631.  Chemical  Properties  of  Carbon. — Any  form  of  carbon 
when  burned  gives  either  carbon  monoxide  or  carbon  dioxide, 
according  to  the  supply  of  oxygen  available. 

If  sulfur  vapor  is  led  over  hot  carbon,  carbon  disulfide,  CS2, 
is  formed.  Carbon  disulfide  is  a  colorless  liquid.  It  is  an 
important  solvent  for  a  variety  of  substances.  An  instance  of 
interest  is  its  use  to  dissolve  rubber.  As  has  already  been 
pointed  out,  it  is  very  easily  inflammable  and  therefore  danger- 
ous (546). 


Carbon  and  Carbon  Compounds  395 

Carbon  does  not  unite  with  hydrogen  (except  to  a  very  small 
extent  at  high  temperatures)  without  the  assistance  of  a  catalyst; 
but  if  hydrogen  is  led  over  a  mixture  of  carbon  and  very  finely 
divided  nickel  at  250°  a  very  good  yield  of  the  gas  methane, 
CH4,  is  obtained. 

Carbon  can  be  made  to  unite  with  metals,  forming  carbides. 
Thus  when  coke  is  heated  with  lime  in  the  electric  furnace 
calcium  carbide  is  formed: 

CaO+3C->CaC2+CO. 

This  substance  is  of  interest  because  of  the  part  it  plays  in  the 
cyanamide  process  for  fixing  atmospheric  nitrogen  (526).  It 
is  also  of  importance  because  of  its  ready  action,  with  water  to 

form  acetylene: 

CaC2+  2H2O  ->  Ca(OH)2+  C2H2.  (49) 

When  sand  is  heated  with  coke  under  special  conditions  car- 
borundum, SiC,  is  formed.  This  important  abrasive  is  ground, 
mixed  with  a  binder,  and  molded  into  grinding  wheels,  knife 
sharpeners,  etc. 

632.  Carbon  Monoxide. — Wherever  combustion  of  carbon 
or  carbon  compounds  goes  on  with  a  deficient  supply  of  air,  the 
odorless,  colorless,  and  very  poisonous  gas  carbon  monoxide  is 
likely  to  be  formed.    It  is  often  seen  burning  as  a  pale-blue 
flame  on  the  top  of  coal  fires.     Care  must  be  taken  that  furnaces 
are  sufficiently  ventilated  to  prevent  this  dangerous  gas  from 
getting  into  living  quarters.     Carbon  monoxide  is  made  in  large 
quantities  for  industrial  fuel  gas.     If  made  from  air  and  coke 
it  is  mixed  with  nitrogen.     Steam  and  coke  give  carbon  monoxide 

and  hydrogen: 

H2O+C->CO+H2. 

633.  Carbon  Dioxide. — Liquid  carbon  dioxide  is  sold  in  steel 
cylinders  in  which  it  is  under  about  60  atmospheres'  pressure. 
When  it  is  allowed  to  flow  rapidly  from  these  into  a  cloth  bag 
it  is  cooled  so  strongly  by  its  own  evaporation  that  tiny  particles 
of  the  solid  carbon  dioxide  are  formed  and  collect  in  the  bag  as 
carbon  dioxide  snow.    This  substance  has  a  vapor  pressure  of 


396  Introduction  to  General  Chemistry 

76  cm.  at  —  79°,  a  temperature  which  is  23°  below  the  melting- 
point  of  the  solid.  Thus  the  latter  evaporates  at  —  79°  without 
melting.  Carbon  dioxide  snow  has  a  limited  but  important  use 
as  a  refrigerant. 

The  gas  is  very  stable  and  will  not  support  ordinary  com- 
bustion. It  is  so  much  heavier  than  air  that  it  can  be  poured 
from  one  vessel  to  another  much  as  we  would  pour  a  liquid. 
These  properties  make  it  very  serviceable  as  a  fire  extinguisher. 
Usually  the  apparatus  contains  concentrated  carbonate  solution 
and  a  bottle  of  sulfuric  acid.  When  the  handle  of  the  device  is 
turned  the  bottle  is  broken,  and  a  concentrated  solution  of 
carbon  dioxide  is  formed  under  considerable  pressure  of  the  gas. 
A  stream  of  the  gas  and  supersaturated  solution  may  then  be 
directed  against  the  fire. 

The  so-called  charged  water  (soda  water)  is,  of  course,  a 
solution  of  carbon  dioxide  under  pressure.  When  the  pressure 
is  released  the  solubility  of  the  gas  decreases,  and  the  liquid  is 
seen  to  froth  with  the  escaping  bubbles. 

Carbon  dioxide  plays  an  important  part  in  plant  life,  as  will 
be  shown  later  (690,  691). 

634.  Illuminating  Gas. — The  distillation  of  coal  yields 
illuminating  gas,  together  with  other  important  products. 
These  vary  in  composition  according  to  the  kind  of  coal  used 
and  the  temperature  to  which  the  latter  is  heated.  The  follow- 
ing data  are  representative  of  the  nature  and  the  approximate 
amounts  of  products  gained  from  one  ton  of  coal : 

Illuminating  gas,  11,000  cu.  ft. 
Ammonia,  6  Ib. 
Coal  tar,  120  Ib. 
Coke,  1,500  Ib. 

The  coal  is  heated  in  retorts,  A,  to  about  1300°  (Fig.  93). 
The  vapors  given  off  are  sent  through  water  in  the  so-called 
hydraulic  main,  B,  in  which  some  of  the  ammonia  and  tar  are 
collected.  The  gas  is  then  passed  first  through  condensers, 
C,  to  take  out  the  rest  of  the  coal  tar,  and  next  through  trays 
of  broken  brick,  D,  " scrubbers,"  over  which  water  is  trickling, 


Carbon  and  Carbon  Compounds 


397 


to  dissolve  the  last  of  the  ammonia.  Next,  objectionable  impu- 
rities such  as  carbon  dioxide  and  hydrogen  sulfide  are  removed. 
Finally  the  gas  is  stored  in  huge  holders  preliminary  to  its 
passage  into  the  city  main.  Illuminating  gas  consists  largely 
of  hydrogen  and  methane,  CH4,  together  with  a  small  proportion 
of  ethylene,  C2H4  (660). 

635.  Organic  Chemistry. — The  number  of  definite  compounds 
containing  the  element  carbon  far  exceeds  that  of  all  other 
chemical  substances.  It  was  once  thought  that  all  except  the 
very  simplest  carbon  compounds  were  products  of  vegetable 
and  animal  organisms,  and  for  this  reason  they  were  called 


FIG.  93 


organic  substances.  We  now  know  that  this  view  is  erroneous, 
because  the  great  majority  of  organic  substances  can  at  present 
be  made  by  purely  chemical  methods  (synthesized)  from  the 
elements  composing  them.  Although  the  term  "organic"  is 
a  misnomer,  it  is  retained  and  serves  to  distinguish  the  great 
class  of  compounds  of  carbon  from  all  others,  which  by  contrast 
are  known  as  inorganic  substances.  A  few  of  the  simplest  com- 
pounds of  carbon  are  classed  as  inorganic;  these  include  such 
substances  as  the  oxides  of  carbon,  carbonates,  carbides  (e.g., 
CaC2),  etc. 

It  will  be  the  purpose  of  the  rest  of  this  chapter  to  give  the 
student  a  brief  glimpse  of  the  fundamental  principles  of  organic 
chemistry,  and  the  object  of  the  next  chapter  to  show  some  of 
the  great  successes  of  this  branch  of  the  science.  No  attempt 
should  be  made  on  the  part  of  the  reader  to  secure  an  intimate 
knowledge  of  the  detail  presented,  since  such  effort  is  best 


398  Introduction  to  General  Chemistry 

expended  under  circumstances  in  which  more  time  is  given  to 
laboratory  work  than  is  possible  in  a  general  chemistry  course. 
But  the  student  should  aim  to  understand  the  fundamental 
principles  pointed  out. 

We  shall  begin  by  the  study  of  one  of  the  commonest  and 
best  known  of  organic  substances,  starch.  We  shall  then  show 
how  from  this  plant  product  a  variety  of  other  substances  can 
be  derived  by  chemical  processes.  This  procedure  will  lead  us 
to  a  knowledge  of  the  properties  and  interrelations  of  several 
of  the  most  important  organic  compounds  and  thus  pave  the 
way  for  the  further  study  of  the  compounds  of  carbon. 

636.  Starch. — Starch  in  nearly  pure  form  is  a  well-known 
article  of  merchandise,  used  extensively  as  food  and  also  for 

TABLE  XXVII 

PERCENTAGE    OF    STARCH   IN    SEVERAL    PLANT 
PRODUCTS 

Wheat 68 

Corn  (maize) 75 

Oats 55 

Rice 78 

Potatoes 15 

Beans 59 

the  stiffening  (starching)  of  laundered  clothes,  etc.  Most  starch 
is  made  from  corn  (maize),  of  which  it  constitutes  about  75  per 
cent.  Starch  forms  a  large  part  of  the  mass  of  most  grains  and 
seeds  and  is  present  in  large  amount  in  tubers,  bulbs,  and  other 
parts  of  plants.  Table  XXVII  shows  the  starch  contents  of  a 
variety  of  plant  products. 

Starch  may  be  easily  prepared  by  grating  a  potato  and 
stirring  the  pulp  with  water.  If  the  mixture  is  poured  through 
a  coffee  strainer,  the  starch  goes  through  with  the  water,  while 
the  fiber  is  left  behind.  The  (insoluble)  starch  soon  settles 
out  of  the  water,  which  can  then  be  decanted.  After  being 
dried  at  room  temperature  the  starch  is  obtained  as  a  white 
powder. 

It  is  well  known  that  starch  is  one  of  the  indispensable  food 
constituents  for  man  and  all  herbivorous  animals.  It  constitutes 


Carbon  and  Carbon  Compounds 


399 


the  most  valuable  ingredient  of  vegetables  and  fruits  and  forms 
a  very  considerable  part  of  all  breadstuffs. 

The  simplest  formula  calculated  from  the  analysis  of  starch 
is  C6HIOOS;  but  since  it  is  not  possible  to  volatilize  this  sub- 
stance its  molecular  weight  cannot  be  found  by  the  vapor- 
density  method.  The  physical  and  chemical  properties  of 
starch  lead  us  to  believe  that  its  formula  is  less  simple  than 
that  given,  and  that  it  is  better  represented  by  (C6HI0Os)n,  where 
n  may  be  a  rather  large  number. 

637.  The  Properties  of  Starch. — Starch  occurs  in  plants  in 
minute  grains  easily  visible  under  the  microscope.     Fig.   94 
shows  the  appearance  of  wheat  starch. 

It  scarcely  need  be  said  that  the  visible 
grains  are  not  the  molecules  of  starch,  for 
each  grain  contains  an  enormous  number 
of  molecules. 

Natural  starch  is  nearly  insoluble  in 
cold  water,  but  when  boiled  with  water 
the    grains    burst,    producing    so-called 
starch  paste.    A  solution  of  starch  does 
not   conduct    an    electric  current  much 
better  than  water;  it  is  not  an  acid,  base,  or  salt.     The  action 
of  iodine  on  starch  (263),  giving  a  blue  color,  forms  the  simplest 
and  best  test  for  this  substance. 

638.  The  Conversion  of  Starch  into  Glucose. — The  most 
important  reaction  of  starch  is  that  in  which  it  unites  with  water 
to  form  glucose,  or  grape  sugar: 

C6HI00S+H20->C6HI206. 

This  reaction  requires  the  aid  of  a  catalytic  agent.  Acids  of  all 
kinds  act  as  catalytic  agents  for  the  conversion  of  starch  into 
glucose.  The  latter  substance  is  made  commercially  by  heating 
starch  with  very  dilute  hydrochloric  acid.  The  acid  rapidly 
promotes  the  union  of  water  with  the  starch  but  is  not  itself 
changed.  At  the  end  of  the  reaction  the  HC1  is  neutralized 
with  soda  and  thus  converted  into  NaCl,  the  presence  of  which 
in  minute  amounts  is  not  objectionable.  The  union  of  starch 


FIG.  94 


4oo  Introduction  to  General  Chemistry 

with  water  is  called  hydrolysis,  and  we  say  that  starch  is  hydro- 
lyzed  to  glucose. 

639.  Glucose,  C6HI2O6. — The  sugar  of  ripe  grapes  and  of 
many  other  fruits  is  largely  glucose,  which  is  also  called  dextrose 
and  grape  sugar.     In  pure  form  glucose  is  a  white  crystalline 
substance  looking  much  like  ordinary  "granulated  sugar. "     It 
is  very  soluble  in  water  and  is  about  half  as  sweet  as  ordinary- 
sugar .     Very  little  pure  glucose  is  made,  but  glucose  sirup  is 
produced  in  enormous  quantities. 

Glucose  sirup  is  used  directly  as  a  table  sirup  and  is  com- 
monly known  as  corn  sirup.  It  is  frequently  colored  and 
flavored  to  imitate  maple  sirup.  It  is  a  good  and  cheap  food- 
stuff. A  large  part  of  the  production  of  glucose  sirup  is  used 
in  the  manufacture  of  candy.  The  sirup  is  now  shipped  largely 
in  tank  cars. 

640.  The  Fermentation  of  Glucose  and  the  Formation  of 
Alcohol. — The  fermentation  of  fruit  juices  is  due  to  the  decom- 
position of  their  glucose  content  into  alcohol,  C2H6O,  and  C02, 
according  to  the  equation 

C6HI206->2C2H60+2CO2. 

This  change  requires  the  aid  of  a  catalytic  agent  produced  by 
growing  yeast,  which  is  a  simple  form  of  vegetable  organism 
(Fig.  95).  It  was  once  thought  that  the 
yeast  consumed  the  glucose  as  food  and 
produced  alcohol  and  C02  as  products. 
The  falsity  of  this  idea  was  shown  by 
grinding  up  growing  yeast  so  as  to 
destroy  every  plant  cell,  then  pressing 
out  the  plant  sap  and  mixing  it  with 
glucose.  The  sap  or  extract  so  obtained 
FIG.  95  fermented  glucose  and  therefore  con- 

tained the  active  catalytic  agent  of  fer- 
mentation. The  change  •  of  glucose  into  alcohol  and  carbon 
dioxide  is  therefore  strictly  a  chemical  decomposition. 

641.  Alcohol,  C2H6O.— There  are  many  kinds  of  alcohols, 
but  the  term  alcohol  is  commonly  and  popularly  applied  to  the 


Carbon  and  Carbon  Compounds  401 

fermentation  product  of  glucose,  C2H6O.  This  is  also  known 
as  grain  alcohol  (because  it  is  usually  made  from  corn,  barley,  or 
rye  as  starting  materials),  and  by  chemists  as  ethyl  alcohol. 
Alcohol  is  a  colorless  liquid  which  burns  with  a  non-luminous, 
sootless  flame.  It  is  miscible  with  water  in  all  proportions. 

In  the  manufacture  of  alcohol,  corn  from  which  the  germ  has 
been  removed  is  ground  up  coarsely  and  boiled  with  water.  The 
product  is  cooled  to  60°  or  65°  and  mixed  with  the  malt,  which 
furnishes  the  diastase  necessary  to  convert  the  starch  into 
maltose,  CI2H22Ou,  which  is  an  intermediate  product  between 
starch  and  glucose.  Maltose  is  easily  changed  into  glucose  by 
heating  with  dilute  HC1  or  H2S04: 


and  like  glucose  is  easily  fermented  to  give  alcohol  and  CO2  : 
CuHaaOxx+HX)  ->4C2H60+4C02. 

After  all  starch  has  changed  into  maltose,  yeast  is  added  to 
start  fermentation,  which  is  finished  in  three  days.  The  product, 
which  contains  10  to  12  per  cent  of  alcohol,  is  then  distilled. 
The  alcohol  (boiling-point  78°)  distils  off  more  readily  than  the 
water  and  is  thus  easily  separated  from  most  of  the  water  as 
well  as  from  the  other  soluble  and  insoluble  materials  present. 
Repeated  distillation  finally  yields  a  product  containing  95  per 
cent  of  alcohol.  Proof  spirit  contains  about  50  per  cent  of 
alcohol.  Pure  alcohol,  free  from  water,  is  called  absolute 
alcohol.  Denatured  alcohol  is  alcohol  to  which  has  been 
added  methyl  alcohol,  benzene,  pyridine,  etc.,  to  render  it  unfit 
for  drinking.  It  is  poisonous. 

642.  Ether,  C4HI00.  —  Ether  is  made  by  allowing  alcohol  to 
drop  slowly  into  a  mixture  of  alcohol  and  two  parts  of  concen- 
trated H2S04  at  140°.  Crude  ether  distils  off  and  is  condensed 
by  cooling  its  vapor  (Fig.  96).  The  reaction  in  its  simplest  form 
may  be  represented  thus: 

2CaH60->C4HIOO+H2O. 


402 


Introduction  to  General  Chemistry 


The  H2S04  acts  as  a  powerful  dehydrating  (water-absorbing) 
agent.  It  is  very  probable  that  the  reaction  is  not  as  simple 
as  indicated,  but  that  an  intermediate  compound  of  CaHeO  and 
H2SO4  is  first  formed  and  later  decomposed. 

Ether  is  a  colorless,  mobile,  and  very  volatile  liquid  boiling 
at  35°.  It  is  our  most  important  anaesthetic,  being  almost 
universally  used  in  surgical  operations.  It  is  also  of  enormous 
importance  chemically.  It  is  an  excellent  solvent  for  numerous 
organic  substances.  It  mixes  with  alcohol  in  all  proportions 
but  is  only  slightly  soluble  in  water. 


FIG.  96 

643.  The  Paraffine  Series. — From  crude  petroleum  and  the 
accompanying  gases  a  very  remarkable  series  of  carbon  and 
hydrogen  compounds,  called  hydrocarbons,  can  be  obtained. 
The  simplest  member  of  the  series  is  methane,  CH4  (54).  It 
was  pointed  out  that  this  gas  is  the  chief  component  of  natural 
gas  and  of  marsh  gas.  Methane  may  be  made  in  the  laboratory 
by  heating  a  mixture  of  sodium  acetate,  sodium  hydroxide,  and 
lime  to  a  high  temperature  in  an  iron  retort.  The  reaction  is 
substantially  as  follows: 

NaC2H302+NaOH  ->  CH4+Na2CO3. 

The  second  member  of  the  series  is  ethane,  C2He,  a  gas  closely 
resembling  methane.  The  third  member  is  propane,  C3Hs. 
The  formula,  name,  physical  state,  boiling-point  (B.P.),  and 
melting-point  (M.P.)  of  a  number  of  these  compounds  are 
given  in  Table  IXXVIII.  It  will  be  seen  that  the  number  of 


Carbon  and  Carbon  Compounds 


403 


H  atoms  is  two  more  than  twice  the  number  of  carbon  atoms  in 
any  molecule.  Therefore  the  general  formula  for  a  compound 
having  n  carbon  atoms  is  CwH2W+2.  All  the  members  of  this 
series  of  compounds  up  to  C20  are  known,  and  also  several  with 
more  carbon  atoms  in  the  molecule. 

TABLE  XXVIII 


B.P. 

B.P. 

M.P. 

CH4  methane,  gas 

-164° 

CvHifi,  heptane,  liquid 

08° 

C2H6  ethane,  gas    .  . 

-  93° 

CgHrg,  octane,  liquid. 

124° 

C3Hg  propane,  gas  

-  ?7° 

CieH34,  liquid  .... 

288° 

1  8° 

C4Hio,  butane,  gas        .    . 

4-    i° 

C20H42,  solid  ... 

77° 

C5Hi2,  pentane,  liquid  

38° 

C24HSo,  solid  

5° 

C6Hi4,  hexane,  liquid  

7i° 

C32H66,  solid  

68° 

The  petroleum  products  in  common  use  are  mixtures  of 
hydrocarbons.  The  principal  components  of  some  of  these  are 
given  in  Table  XXIX. 

TABLE  XXIX 

Gasoline C5HI2  to  C8Hl8 

Kerosene C8Hl8  to  CI2H26 

Lubricating  oils CI2H26  to  CZ6H34 

Vaseline C^H^  to  C20H42 

Paraffine C24HSO  to  C36H74 

644.  The  Action  of  Chlorine  on  Methane. — Chlorine  acts  on 
methane  to  form  a  series  of  four  compounds.  The  first  reaction 
yields  methyl  chloride,  CH3C1: 

CH4+C12->CH3C1+HC1. 

Methyl  chloride  is  a  colorless  gas  which  can  react  still  further 

with  chlorine : 

CH3C1+  C12  ->  CH2C12+HC1. 

The  product,  CH2C12,  methylene  chloride,  is  a  colorless 
liquid,  which  gives,  with  more  chlorine,  chloroform,  CHC13: 

CH2C13+ Cla  ->  CHC13+HC1. 

Chloroform  is  a  sweet-smelling,  colorless  liquid  boiling  at  61° 
and  insoluble  in  water,  but  miscible  with  alcohol  or  ether  in  all 


404  Introduction  to  General  Chemistry 

proportions.  It  is  a  good  anaesthetic  but  is  not  as  safe  as  ether 
for  patients  with  weak  hearts.  Finally  the  action  of  more  C12 
on  CHC13  gives  carbon  tetrachloride  : 

CHC13+C12->CC14+HC1. 

This  is  a  colorless  liquid  which  is  not  only  incombustible  but 
an  excellent  fire  extinguisher.  It  is  the  principal  component 
used  in  Pyrene  and  similar  fire  extinguishers.  Devices  of  this 
type  should  be  at  hand  in  every  laboratory  for  the  speedy  control 
of  fires  that  might  otherwise  prove  dangerous. 

Carbon  tetrachloride  and  closely  related  substances  are  used 
as  cleaning  fluids.  They  are  perfectly  safe,  since  they  are  not 
combustible. 

645.  Methyl  Alcohol,  CH4O.—  If  methyl  chloride,  CH3C1,  is 
heated  with  water  at  a  high  temperature  under  pressure  it  reacts 
as  follows: 

CH3C1+H2O  ->  CH40+HC1. 

The  substance,  CH40,  is  methyl  alcohol.  This  method  of  mak- 
ing methyl  alcohol  is  only  of  scientific  interest.  The  substance 
is  obtained  in  large  quantity  as  one  of  the  volatile  products 
of  the  distillation  of  wood  (701).  The  common  name  of  methyl 
alcohol  is  wood  alcohol.  It  is  a  colorless  liquid,  boiling  at  66° 
and  miscible  with  water  in  all  proportions  It  resembles  com- 
mon (ethyl)  alcohol  in  its  physical  and  chemical  properties. 
It  is  a  dangerous  poison,  often  accidentally  causing  permanent 
blindness  or  death.  It  is  a  valuable  solvent  and  is  manufactured 
on  a  large  scale.  « 

646.  Methyl  Ether,  C2H6O.—  Methyl  ether  is  made  from 
methyl  alcohol  in  practically  the  same  way  that  ethyl  ether, 
C4HIDO,  is  made  from  ethyl  alcohol: 


Methyl  ether,  C2H60,  is  a  colorless  gas  having  chemical  prop- 
erties resembling  those  of  ethyl  ether. 

647.  Isomerism.  —  The  formula  given  for  methyl  ether  is 
that  based  upon  analysis  and  gas  density.     It  will  be  recalled 


Carbon  and  Carbon  Compounds  405 

that  exactly  the  same  formula,  C2H60,  was  also  ascribed  to  ethyl 
alcohol.  This  formula  was  also  fixed  by  analysis  and  vapor 
density.  Both  substances,  although  entirely  distinct  in  physical 
and  chemical  properties,  have  exactly  the  same  percentage  com- 
position, the  same  vapor  density,  and  therefore  the  same  molecular 
weight!  This  is  a  most  remarkable  fact  but  by  no  means  the 
only  known  case  of  the  kind.  Indeed,  among  organic  com- 
pounds there  are  hundreds,  yes  thousands,  of  cases  in  which 
two  or  even  several  totally  different  substances  have  the  same 
percentage  composition  and  the  same  molecular  weight  and 
in  consequence  the  same  formula.  Such  pairs  or  groups  of 
substances  are  called  isomers.  In  consequence  we  speak  of 
the  isomerism  of  methyl  ether  and  ethyl  alcohol.  We  shall 
next  consider  some  theoretical  matters  with  the  object  of 
developing  an  explanation  of  the  remarkable  phenomenon  of 
isomerism. 

648.  The  Valence  of  Carbon. — If  we  consider  the  six  com- 
pounds of  carbon,  CH4,  C2H2,  C2H4,  C2H6,  C3H6,  and  C3H8  (see 
Table  IV,  63),  we  shall  have  much  difficulty  in  deciding  upon 
the  valence  of  carbon.  In  CH4,  assuming  the  valence  of  hydro- 
gen to  be  one,  carbon  evidently  has  a  valence  of  four,  but  this 
value  does  not  seem  to  harmonize  with  the  formulae  of  the  other 
substances.  Consistent  conclusions  are  to  be  reached  only  by 
considering  the  question  of  the  graphic  formulae  of  the  substances 
(323).  If  we  start  with  the  plausible  assumption  that  carbon 
has  a  valence  of  four,  while  hydrogen  and  chlorine  each  have 
unit  valence,  we  may  write,  as  the  graphic  formulae  of  methane 
and  the  substances  formed  from  it  by  the  action  of  chlorine,  the 
following : 

H  H  H  H  Cl 

H— C— H       H— C— Cl       H— C— Cl       Cl— C— Cl  Cl— C— Cl 

H  H  Cl  Cl  Cl 

Since  we  write  for  water  the  graphic  formula 

H-O-H,  (323) 


406  Introduction  to  General  Chemistry 

thereby  assuming  the  valence  of  oxygen  to  be  two,  we  may 
write  for  methyl  alcohol,  CH4O  (645),  the  graphic  formula 

H 
H—  C—  O—  H 


It  seems  probable  that  these  graphic  formulae  show  the  actual 
relations  of  the  atoms  to  one  another  in  the  molecule.  If  so, 
they  show  the  structure  of  the  molecule,  and  in  consequence 
they  may  be  called  structural  formulae.  As  stated  earlier  (323), 
the  lines  joining  the  symbols  in  such  formulae  are  called  bonds. 
Only  one  possible  arrangement  of  two  atoms  of  carbon  and 
six  atoms  of  hydrogen  (C2H6)  satisfies  the  condition  that  the 
valence  of  carbon  is  four  and  that  of  hydrogen  one,  namely  the 

following  : 

H    H 

H—  C—  C—  H 

I       i 
H    H 

This  is  therefore  the  accepted  formula  of  ethane  (643). 

649.  The  Structural  Formulae  of  Ethyl  Alcohol  and  Methyl 
Ether.  —  Ethyl  alcohol  and  methyl  ether  are  both  represented 
by  the  simple  formula  C2H60.  There  are  two  structural 
possibilities! 

H    H  H          H 

II  I  ! 

H—  C—  C—  O—  H  and  H—  C—  O—  C—  H 

I  I 

H    H  H          H 

Since  methyl  alcohol  is 

H 

H—  C—  O—  H 


and  since  methyl  and  ethyl  alcohol  show  similar  chemical 
behavior,  we  may  assume  provisionally  that  the  first  of  the 


Carbon  and  Carbon  Compounds  407 

preceding  formulae  is  that  of  ethyl  alcohol  and  the  second  that 
of  methyl  ether. 

The  formation  of  ethyl  ether  from  ethyl  alcohol  could  then 
be  considered  to  take  place  by  the  removal  of  a  molecule  of 
water  from  each  pair  of  alcohol  molecules,  thus: 

H    H  H   H 


4-C— C— H 

J   |      | 
H    H  H   H 


H— C— C—O-j-H+H— OH 


which  leads  to  the  following  as  the  formula  for  ethyl  ether, 
C4HI00: 

H    H          H    H 

I     -        I      I 
H— C— C— O— C— C— H 

I  I      I 

H    H  H    H 

This  conclusion  is  supported  by  the  fact  that  this  formula  is 
entirely  analogous  to  that  assigned  to  methyl  ether.  In  general, 
the  structure  of  the  molecules  of  any  substance  is  discovered 
by  a  study  of  its  reactions  and  mode  of  synthesis. 

650.  The  Cause  of  Isomerism. — Isomerism  finds  a  simple 
explanation  in  the  assumption  that  the  atoms  composing  mole- 
cules of  isomers  are  arranged  in  two  (or  more)  ways,  each  in 
harmony  with   simple  valence  laws.     Hundreds  of  cases   of 
isomerism  are  all  adequately  explained  in  this  fashion. 

651.  Discussion. — We  have  now  seen  how  a  number  of  the 
simpler  organic  compounds  can  be  obtained  from  natural  sources. 
We  have  also  learned  a  little  regarding  the  reactions  of  organic 
substances. 

The  disclosure  of  the  isomerism  of  ethyl  alcohol  and  methyl 
ether  demanded  an  explanation  which  was  found  in  the  assump- 
tion that  the  atoms  of  a  molecule  are  joined  to  one  another  in 
definite  fashion  and  always  in  accord  with  simple  fixed  rules  of 
valence. 

The  importance  of  structural  formulae  to  the  chemist  can 
well  be  illustrated  with  the  examples  of  methyl  ether  and  ethyl 


408  Introduction  to  General  Chemistry 

alcohol.  The  OH  group  attached  to  a  carbon  atom  which  carries 
no  other  oxygen  atom  is  called  an  alcohol  group.  When  an 
organic  chemist  sees  this  group  in  a  structure  he  knows  that  the 
substance  must  have  the  properties  of  an  alcohol  modified  to 
some  degree,  according  to  what  other  groups  are  near  it  in  the 
molecule  in  question.  An  oxygen  atom  joined  on  either  hand 
to  a  carbon  atom  which  carries  no  other  oxygen  atoms  is  the 
group  characteristic  of  the  oxygen  ethers.  The  latter  are  very 
different  substances  from  alcohols  and  have  their  own  character- 
istics. Thus  at  a  glance  the  trained  chemist  can  read  the  char- 
acteristics of  a  substance  by  observing  the  arrangements  of  the 
atoms  in  the  structure  formula  of  the  molecule  and  recognizing 
groups  which  have  pronounced  characteristics. 

Most  of  the  substances  encountered  in  preceding  chapters 
belong  to  one  or  another  of  five  principal  classes :  elements,  oxides, 
acids,  bases,  or  salts.  Among  organic  compounds  we  find  a  large 
number  of  new  and  important  classes.  The  paraffine  hydro- 
carbons, CnH2»+2,  starting  with  methane  form  the  simplest  class 
of  organic  substances.  The  halogen  derivatives  like .  methyl 
chloride,  CH3C1,  and  chloroform,  CHC13,  form  another  class. 
The  alcohols  and  the  ethers  are  also  important  classes.  The 
balance  of  this  chapter  will  be  devoted  to  the  brief  description 
of  several  other  classes  of  carbon  compounds.  It  is  necessary 
that  the  student  should  have  at  least  a  slight  acquaintance 
with  the  more  important  classes  of  organic  substances  if  he 
wishes  to  get  an  insight  into  the  nature  of  foodstuffs,  as  well  as 
of  those  organic  substances  which  play  so  important  a  role  in 
our  modern  daily  life.  Among  the  latter  are  found  medicinals, 
dyes,  explosives,  photographic  developers,  perfumes,  poisons,  etc. 

It  is  not  to  be  expected  that  the  student  who  reads,  however 
carefully,  the  balance  of  this  chapter  will  get  very  definite  ideas 
of  the  methods  of  preparation  and  properties  of  the  various 
classes  of  substances  there  described.  Familiarity  with  organic 
substances  can  be  gained  only  by  prolonged  and  detailed  study 
in  both  text  and  laboratory.  Our  object  in  presenting  the  topics 
about  to  be  discussed  will  be  largely  attained  if  we  impress  upon 
the  student  that  carbon  forms  an  enormous  variety  of  com- 


Carbon  and  Carbon  Compounds  409 

pounds,  that  these  compounds  are  of  known  molecular  structure, 
that  they  fall  into  definite  classes  according  to  the  presence  in 
them  of  certain  active  groups  of  atoms,  and  that  the  members 
of  each  class  show  similar  behavior,  namely,  that  of  the  active 
groups,  the  characteristics  of  each  of  which  may  be  modified 
somewhat  according  to  the  nature  of  other  groups  present  in 
the  same  molecule. 

In  what  follows,  bonds  between  characteristic  groups  will 
frequently  be  represented  by  dots  or  entirely  omitted.  Thus 
CH3  —  OH  will  be  written  CH3  •  OH,  or  for  still  greater  simplicity 
CH3OH. 

652.  Aldehydes.  —  The  mild  oxidation  of  ethyl  alcohol  gives 
acetalde'hyde.  The  reaction  is  best  carried  out  with  a  mixture 
of  a  dichromate  and  sulfuric  acid.  Omitting  details,  the  equa- 

tion is 

CH3  •  CH2  -  OH+O^CH3  •  CO  -  H+H2O. 


Acetaldehyde  is  a  very  volatile  liquid. 

The  oxidation  of  methyl  alcohol,  CH3OH,  gives  the  closely 
related  substance  formaldehyde,  H  •  CO  •  H.  This  is  a  gas  of 
pungent  odor  which  comes  on  the  market  as  a  40  per  cent  solu- 
tion in  water  called  formalin.  It  is  extensively  used  as  a 
germicide  and  antiseptic  and  also  in  the  manufacture  of  other 
important  organic  substances.  In  general,  the  oxidation  of  any 
alcohol  of  the  formula  R  •  CH2  •  OH,  where  R  is  H  or  a  hydro- 
carbon radical,  as  for  example  C2HS  or  C3H7,  gives  an  aldehyde, 
R  •  CO  •  H.  The  aldehydes  are  good  reducing  agents  because 
they  are  readily  oxidized  to  the  corresponding  acids.  For 
example,  acetaldehyde  may  be  oxidized  by  suitable  agents  to 

acetic  acid: 

CH3  •  CO  -  H+0->CH3  •  CO  •  OH. 

653.  Acetic  Acid,  HC2H3O2.  —  As  stated  earlier  (157),  acetic 
acid  is  the  most  important  component  of  vinegar,  of  which  it 
forms  about  4  per  cent.  Cider  vinegar  is  made  from  cider,  the 
juice  of  apples.  The  sweetness  of  fresh  cider  is  due  to  glucose. 
Upon  standing  several  days  the  glucose  gradually  ferments, 
owing  to  the  growth  of  yeast,  the  germs  of  which  are  always 


410  Introduction  to  General  Chemistry 

present  in  the  dust  of  the  air,  the  glucose  changing  to  alcohol  and 
C02.  The  fermented  product  is  popularly  known  as  hard  cider. 
If  hard  cider  is  allowed  to  stand  in  an  open  or  loosely  stoppered 
vessel  it  changes  in  a  few  weeks  into  vinegar.  This  change  is 
the  result  of  the  oxidation  of  alcohol  to  acetic  acid : 

C2H60+02  -»HC2H302+H20. 

The  oxidation  requires  the  assistance  of  a  catalytic  agent  pro- 
duced by  the  so-called  vinegar  plant,  or  mother  of  vinegar 
(micoderma  aceti).  "White  distilled  vinegar"  is  made  from 
dilute  alcohol  produced  from  corn,  substantially  in  the  manner 
already  described  (640).  The  dilute  alcohol  is  allowed  to  trickle 
slowly  through  large  casks  filled  with  beech  wood  shavings, 
coated  with  the  slimy  mother  of  vinegar,  while  oxygen  is 
furnished  by  a  countercurrent  of  air  that  enters  near  the  bot- 
tom and  passes  out  at  the  top  of  the  cask.  This  so-called 
"  quick  process"  produces  finished  vinegar  in  eight  to  ten  days. 

Pure  acetic  acid  is  a  colorless  liquid  having  a  sharp,  char- 
acteristic odor.  When  free  from  water  it  solidifies  at  17°  to 
glassy  crystals  called  glacial  acetic  acid.  We  have  already 
studied  the  reactions  of  solutions  of  acetic  acid  and  its  salts 
(157,424,456). 

654.  The  Graphic  Formula  of  Acetic  Acid. — Expressed  graphi- 
cally, we  have  as  the  equation  for  the  oxidation  of  alcohol, 

H    H  H 

I  I  I 

H— C— C— O— H+O2->H— C— C— O— H+HaO 

II  I      II 
H    H                              HO 

The  group   -C-OH,  or  briefly   -CO  •  OH,  is  the  carboxyl 
O 

radical.  One  of  the  oxygen  atoms  of  acetic  acid  is  said  to  be 
attached  to  one  of  the  carbon  atoms  by  a  double  bond  (324). 
Of  the  four  H  atoms  of.  acetic  acid,  one  occupies  a  unique  posi- 
tion hi  that  it  is  attached  to  an  atom  of  O,  while  the  others  are 
attached  to  one  of  the  carbon  atoms.  Since  only  one  of  the  H 


Carbon  and  Carbon  Compounds  411 

atoms  of  the  molecule  is  ionizable  we  may  safely  conclude  that 
it  is  the  one  attached  to  oxygen. 

655.  The  Fatty  Acids. — There  is  a  series  of  acids  closely 
related  to  acetic  acid,  all  having  a  carboxyl  radical  attached 
to  a  hydrocarbon  radical.     The  general  formula  of  such  acids 
is  R  •  CO  •  OH,  where  R  stands  for  hydrogen  or  any  hydro- 
carbon radical  such  as  C2HS,  C3H7,  C4H9,  etc.    The  names  and 
formulae  of  a  few  of  the  more  important  are  given  in  Table  XXX. 

TABLE  XXX 

Formic  acid H  •  CO  •  OH. 

Acetic  acid CH3  -  CO  •  OH 

Propionic  acid C2HS  •  CO  •  OH 

Butyric  acid C3H7  •  CO  •  OH 

Palmitic  acid CISH,X  •  CO  •  OH 

Stearic  acid CI7H3S  •  CO  •  OH 

These  acids  are  known  as  fatty  acids,  because  some  of  them  are 
obtained  from  fats  (680). 

656.  Ketones. — Calcium  acetate,  Ca(C2H302)2,  is  an  impor- 
tant article  of  commerce  made  by  the  action  of  acetic  acid  on 
limestone  (CaCO3).    When  calcium  acetate  is  strongly  heated 
it  decomposes  into  CaC03  and  acetone,  C3H6O: 

Ca(C2H3O2)2  ->  CaCO3+ C3H6O. 

The  acetone  distils  off  and  is  condensed  to  a  liquid.  Purified 
acetone  is  a  colorless  liquid  of  mild  but  peculiar  odor.  It  boils 
at  56°.  It  mixes  with  water  or  alcohol  in  all  proportions.  It  is 
an  excellent  solvent  for  many  organic  substances,  and  it  is  also 
used  in  the  preparation  of  several  important  organic  compounds, 
of  which  chloroform  is  one.  The  structural  formula  of  acetone 
is  indicated  by  its  formation  from  calcium  acetate: 


>c=o 


412  Introduction  to  General  Chemistry 

Acetone  is  the  simplest  member  of  a  class  of  substances  called 
ketones.  The  general  formula  for  a  ketone  is  Rz  •  CO  •  R2, 
where  Rr  and  R2  represent  the  formulae  of  hydrocarbon  radicals. 
In  acetone  Rj  and  R2  are  both  methyl,  CH3;  but  if,  say,  Rr  is 
ethyl,  C2HS,  and  R2  is  propyl,  C3H7,  the  formula  of  the  ketone 
would  be  C2HS  •  CO  •  C3H7. 

Ketones  are  closely  related  to  aldehydes,  since  if  R2  is  H  we 
have  RI  •  CO  •  H,  the  formula  of  an  aldehyde.  They  are 
reducing  agents  but  are  not  as  active  as  aldehydes. 

657.  Esters. — The  esters  are  an  important  class  of  com- 
pounds, inasmuch  as  all  animal  and  vegetable  fats  and  oils  are 
included  therein.  One  of  the  simplest  esters,  ethyl  acetate, 
CH3COOC2H5,  is  obtained  by  the  action  of  ethyl  alcohol  on 
acetic  acid: 

CH3  •  CO  •  OH+C2HS  -  OH->CH3  •  CO  -  OC2HS+H2O. 

Ethyl  acetate  is  a  colorless  liquid,  boiling  at  75°.  It  has  a 
rather  pleasant  odor  and  is  somewhat  soluble  in  water. 

Just  as  most  acids  can  form  salts  with  most  bases,  so  most 
acids  can  form  esters  with  most  alcohols.  However,  alcohols 
are  not  bases,  since  they  do  not  yield  OH~  ions,  and  esters  are 
not  salts.  Their  water  solutions  are  not  ionized,  and  they  do 
not  give  the  ionic  reactions  shown  by  solutions  of  the  acids  from 
which  they  are  derived.  The  formation  of  esters  requires  in 
general  the  stimulus  of  H+  ions  as  a  catalytic  agent.  In  mak- 
ing ethyl  acetate  we  add  to  the  mixture  of  acetic  acid  and  alcohol 
some  HC1  or  H2S04.  The  union  of  acetic  acid  and  alcohol  does 
not  take  place  completely  but  reaches  equilibrium  when  about 
two- thirds  of  the  possible  amount  of  ester  has  been  formed. 
This  is  because  the  reaction  is  reversible. 

The  change  of  ester  and  water  into  acid  and  alcohol  is  pro- 
moted by  the  presence  of  much  water  and  also  by  the  catalytic 
influence  of  acids.  The  speed  of  ester  formation  and  also  the 
speed  of  reaction  of  ester  and  water  are  greatly  increased  with 
increase  of  temperature. 

Methyl  acetate,  CH3CO  •  OCH3,  closely  resembles  ethyl 
acetate.  It  is  miscible  with  water  in  all  proportions.  Other 


Carbon  and  Carbon  Compounds  413 

esters,  including  fats  and  oils,  will  be  considered  in  the  next 
chapter. 

658.  Amines.  —  The  amines  are  derivatives  of  ammonia  and, 
like  the  latter,  are  base-forming  substances  capable  of  yielding 
salts  with  acids.  The  simplest  member  of  the  class  is  methyl 
amine,  CH3NH2.  If  methyl  iodide  (660)  and  ammonia  are 
mixed  they  unite  to  form  methyl  ammonium  iodide: 

CH3I+NH3->CH3NH3I. 

The  product  is  a  soluble  salt  of  the  base  CH3NH3OH.  This 
unstable  base,  which  is  set  free  by  the  action  of  sodium  hydroxide 
on  the  salt,  easily  dissociates  into  methyl  amine  and  water. 
Methyl  amine  is  a  colorless  gas  with  an  odor  resembling  ammonia. 
It  is  abundantly  soluble  in  water,  with  which  it  partially  unites, 

thus: 

CH3  -  NH2+H2O^CH3  -  NH3  •  OH. 


Ethyl  iodide,  C2HSI,  and  ammonia  give  ethyl  ammonium 
iodide,  C2HS  •  NH3  •  I,  from  which  we  readily  obtain  ethyl 
amine,  C2H5  •  NH2,  a  substance  closely  resembling  methyl 
amine. 

Methyl  amine  acts  on  methyl  iodide  as  follows: 

CH3  -  NHa+CH3I->(CH3)aNHaI, 

from  which  dimethyl  amine,  (CH3)2NH,  is  obtained  by  the 
action  of  alkali.  Dimethyl  amine,  by  further  action  of  methyl 
iodide,  yields  (CH3)3NHI,  from  which  by  the  action  of  alkali 
we  get  trimethyl  amine,  (CH3)3N  (59).  Amines  of  various  kinds 
are  usually  found  among  the  products  of  decomposition  of 
proteins  (685).  Trimethyl  amine,  for  example,  is  contained  in 
herring  brine. 

659.  Amides.  —  The  interaction  of  ethyl  acetate  and  ammonia 
gives  acetamide  and  alcohol  : 

CH3COOC2HS+NH3  ->  CH3CONH2+  C2HSOH. 

Acetamide  is  also  easily  made  by  distilling  ammonium  acetate: 
CH3COONH4  ->€H3CONH2+H20. 


414  Introduction  to  General  Chemistry 

The  substance  is  a  white  crystalline  solid  easily  soluble  in  water. 
It  unites  with  hydrochloric  acid  to  form  a  saltlike  compound, 
CH3CONH3C1.  This  fact  shows  that  the  NH2  radical  in  an 
amide  has  still  some  basic  properties. 

Most  organic  acids  and  some  mineral  acids  are  able  to  form 
amides.  Carbonic  acid,  for  example,  forms  the  amide  CO(NH2)2. 
This  important  substance  is  commonly  known  as  urea.  By  far 
the  larger  part  of  the  nitrogen  content  of  the  food  of  all  animals 
is  excreted  as  urea.  Urea  reacts  slowly  with  water  to  form 
ammonia  and  carbon  dioxide: 

CO(NH2)2-hH2O  ->  2NH3+  C02. 

660.  Ethylene,  C2H4. — A  modification  of  the  process  of  mak- 
ing ether  (642)  yields  ethylene.  To  make  ethylene,  a  mixture 
of  alcohol  with  six  parts  by  weight  of  concentrated  sulfuric  acid 
is  heated  to  165°,  and  a  mixture  of  one  part  of  alcohol  to  two 
parts  of  sulfuric  acid  is  dropped  in  slowly.  The  gas  C2H4  is 
given  off.  The  reactions  in  this  case  are  probably  also  complex, 
but  the  net  result  is  the  decomposition  of  alcohol  into  water  and 
ethylene : 

C2H6O->H2O+C2H4. 

Ethylene  is  a  colorless  gas  nearly  insoluble  in  water.  It  burns 
with  a  luminous  flame.  It  gives  several  interesting  and  impor- 
tant reactions.  With  C12  it  unites  to  form  ethylene  chloride: 

C2H4+C12->C2H4C12, 

a  colorless,  oil-like  liquid,  boiling  at  84°.  Ethylene  chloride  is 
insoluble  in  water.  It  has  none  of  the  properties  of  a  salt. 
Ethylene  also  unites  with  bromine,  thus: 

C2H4+Br2->C2H4Br2. 

The  product,  ethylene  bromide,  resembles  the  chloride.  Eth- 
ylene and  HBr  unite  readily  to  form  C2HsBr,  ethyl  bromide,  a 
liquid  boiling  at  38°: 

C2H4+HBr-»C2H5Br. 


Carbon  and  Carbon  Compounds  415 

Ethylene  and  HI  give  ethyl  iodide,  boiling-point  72°: 
C3H4+HI->C2HSI. 

Neither  ethyl  bromide  or  ethyl  iodide  has  any  of  the  properties 
of  a  salt.  Methyl  iodide,  CH3I,  closely  resembles  ethyl  iodide. 
It  is  used  in  the  preparation  of  other  organic  compounds  (658). 
661.  The  Structural  Formula  of  Ethylene.— The  fact  that 
ethylene  unites  with  chlorine,  hydrobromic  acid,  etc.,  leads  to 
the  conclusion  that  two  of  the  valence  bonds  of  the  carbon 
atoms  of  the  C2H4  molecule  are  either  free, 


H    H 
H- 


xx        -ii 

[— C— C— H, 


or,  more  probably,  are  attached  to,  or  satisfied  by,  one  another 
so  as  to  form  a  double  bond  (324)  between  the  two  carbon  atoms  : 

H    H 

I       I 
H—  C=C—  H. 

In  the  reaction  with  chlorine,  for  example,  the  extra  bonds  unite 

with  chlorine  to  give 

H    H 

H—  C—  C—  H. 
Cl   Cl 

662.  The  Ethylene  Series,  CnH2M.  —  A  long  series  of  hydro- 
carbons which  have  twice  as  many  H  as  C  atoms  per  molecule 
and  of  which  the  first  member  is  C2H4  is  known.  The  second 
member  is  propylene,  CH3  —  CH=  CH2.  There  are  theoretically 
just  three  butylenes,  C4H8  : 


(1)  CH3  •  CH2 

(2)  CH3  -  CH  =  CH  •  CH3 


(3) 


4i 6  Introduction  to  General  Chemistry 

These  are  all  known.    All  of  the  members  of  this  series  unite 
with  C12,  Br2,  HI,  etc.     For  example, 

CH3  -  CH2  -  CH  =  CH2-f-  C12  ->  CH3  -  CH2  -  CHC1  -  CH2C1. 

663.  The  Acetylene  Series,  CWH2«_2. — The  acetylene  series 
of  hydrocarbons  has  the  general  formula  CWH2M_2.  It  is  headed 
by  acetylene,  C2H2  (49,  83).  It  is  probable  that  the  four  extra 
valence  bonds  of  acetylene, 

H— C— C— H, 


are  united  in  pairs,  so  that  the  carbon  atoms  are  joined  by  a 
triple  bond, 

H— CEEEC— H. 

In  accord  with  this  view  we  should  expect  acetylene  to  unite 
with  chlorine,  thus: 

Cl    Cl 

H— C=C— H+  2C12-H— C— C— H 

Cl    Cl 

and,  in  fact,  it  actually  behaves  in  this  way.  Because  of  the 
fact  that  C2H4  and  C2H2  have  unsatisfied  valences,  as  shown 
by  their  union  with  C12,  HBr,  etc.,  these  hydrocarbons  are  said 
to  be  unsaturated. 

664.  Isomerism  of  Hydrocarbons. — All  of  the  compounds 
CwH2n+2,  where  n  is  four  or  more,  can  exist  in  two  or  more 
isomeric  forms.  The  simplest  case  is  that  of  the  butanes, 
C4H10,  of  which  two  are  known:  normal  butane, 

CH3-CH2-CH2-CH3, 

and  isobutane, 

CH3v 

>CH-CH3. 
Ctt/ 


Carbon  and  Carbon  Compounds  417 

There  are  three  isomeric  pentanes: 

(1)  CH3  -  CH2  •  CH2  •  CH2  •  CH3 

CH3v 

(2)  >CH-CH2-CH3 
CH/ 

CH3 

(3)  CH3-C-CH3 

CH3 

The  first  of  these  is  said  to  have  a  straight  carbon  chain;  the 
second  and  third  have  branched  chains.  The  B.P.  and  M.P. 
columns  of  Table  XXVIII  (643)  refer  to  the  normal  hydro- 
carbons, with  straight  chains. 

665.  Some  Common  Organic  Acids.— We  shall  now  briefly 
describe  several  of  the  commoner  organic  acids.  The  simplest 
fatty  acid  is  formic  acid,  HCOOH.  It  is  present  in  the  bodies 
of  ants  and  constitutes  the  poison  of  bees'  stings.  It  can  be 
made  artificially  in  several  ways.  It  is  a  colorless  liquid  dis- 
solving easily  in  water  to  form  a  moderately  well-ionized  acid 
solution. 

Oxalic  acid,  H2C2O4,  is  a  white  crystalline  substance.     Its 

structural  formula  is 

COOH 


COOH. 


It  is  a  dibasic  acid  (102),  forming  both  acid  and  neutral  salts, 
as  for  example  NaHC2O4  and  Na2C204.  The  free  acid,  which  is 
easily  soluble  in  water  and  rather  highly  ionized,  crystallizes 
from  water  as  a  hydrate  of  the  formula  H2C204  •  2H20.  The 
acid  is  decidedly  poisonous,  probably  because  of  the  ease  with 
which  it  decomposes  into  the  powerful  poison  carbon  monoxide 

(632)  and  water: 

H2C2O4->H2O+2CO. 

This  decomposition  takes  place  rapidly  when  the  acid  is  heated 
with  concentrated  sulfuric  acid,  and  so  affords  a  good  way  of 
making  carbon  monoxide.  Oxalic  acid  occurs  in  some  plants. 


4i 8  Introduction  to  General  Chemistry 

The  chief  acid  present  in  sour  milk  is  lactic  acid,  the  graphic 
formula  of  which  is  CH3  •  CHOH  -  COOH.  It  gives  a  colorless 
water  solution  of  a  pleasant  sour  taste.  Tartaric  acid,  a  dibasic 
acid  having  the  formula  H2C4H4O6,  is  abundant  in  grapes  in  the 
form  of  its  acid  potassium  salt,  KHC4H4O6.  This  salt  is  obtained 
as  argol  in  large  amount  in  the  manufacture  of  wine.  The 
refined  salt  is  known  as  cream  of  tartar.  High-grade  baking 
powder  is  a  mixture  of  cream  of  tartar  and  sodium  bicarbonate. 
The  dry  mixture  is  fairly  stable.  In  the  presence  of  water  the 
substances  react  thus: 

KHC4H4O6+NaHCO3->KNaC4H406+CO2+H2O. 

The  leavening  power  of  baking  powder  is  due  to  the  CO2  given 
off  (cf.  593).  Tartaric  acid  forms  colorless  crystals  easily  soluble 
in  water.  Its  graphic  formula  is 

0 


COOH                           C  -  OH 

CHOH                     H  •  ( 

:  •  OH 

1                    or 

CHOH                     H  •  ( 

1 

:  -OH 

COOH  C  •  OH 

i! 
o 

Citric  acid,  H3  •  C6HS07,  is  the  acid  of  lemon  and  other 
citrus  fruits.  It  is  a  tribasic  acid  (159)  rather  closely  resembling 
tartaric  acid  but  having  a  more  complex  formula.  Prussic  or 
hydrocyanic  acid,  HNC,  is  an  extremely  poisonous  and  very 
weak  acid.  Its  salts,  sodium  cyanide,  NaNC,  and  potassium 
cyanide,  KNC,  are  made  from  calcium  cyanamid  (526).  Both 
salts  react  readily  with  either  oxygen  or  sulfur,  forming  the 
corresponding  cyanate  or  sulfocyanate  respectively: 

2KNC+O2-»2KNCO, 
KNC+S->KNCS. 

The  cyanide  ion  unites  with  many  metal  ions  to  form  com- 
plex cyanide  ions.  Salts  containing  these  ions,  such  as  potas- 


Carbon  and  Caroon  Compounds  419 

slum  ferrocyanide,   K4Fe(CN)6,   and  potassium  ferricyanide, 
K3Fe(CN)6,  are  important  analytical  reagents. 

666.  The  Aliphatic  and  Aromatic  Series  of  Organic  Com- 
pounds.— Broadly  speaking,  organic  compounds  constitute  two 
great  series,  the  aliphatic  and  the  aromatic.  The  substances 
thus  far  mentioned  are  all  members  of  the  aliphatic  series.  We 
may  consider  that  the  paraffin  hydrocarbons,  CnH2n+2,  are 
the  fundamental  substances  from  which  all  aliphatic  compounds 
are  derived.  Thus  paraffin  hydrocarbons  by  loss  of  two 
hydrogen  atoms  per  molecule  give  members  of  the  ethylene 
series,  CMH2n,  or  by  loss  of  four  hydrogen  atoms  per  molecule 

TABLE  XXXI 

Formula  of  Typical 
Class  of  Compound  Compound 

Hydrocarbon CH3     H 


Halide CH3 

Alcohol CH3 

Aldehyde CH3 

Acid CH3 

Ester CH, 

Amide CH3 

Ether CH3 

Ketone CH3 

Amine CH3 


Cl 
OH 

CO-  H 
CO -OH 
CO  •  OCH3 
CO  •  NH2 
O-  CH3 
CO  -  CH3 
NH2 


give  members  of  the  acetylene  series.  Substitution  of  hydrogen 
by  halogens  or  by  various  radicals  such  as  methyl,  CH3; 
hydroxyl,  OH;  carboxyl,  COOH;  amid,  NH2;  etc.,  gives  rise 
to  the  various  classes  of  compounds  already  briefly  studied. 

Table  XXXI  shows  the  formulae  of  typical  members  of  the 
most  important  classes  of  aliphatic  compounds. 

Instead  of  the  methyl  radical,  CH3  (Table  XXXI),  we  may 
have  ethyl,  C2HS,  propyl,  C3H7,  or  any  radical  derived  from  any 
hydrocarbon  in  each  case  and  thus  obtain  a  very  great  variety 
of  substances.  All  of  these  substances  belong  to  the  aliphatic 
series. 

The  aromatic  compounds  differ  from  the  aliphatic  in  that 
they  are  derived  from  hydrocarbons  entirely  different  from  the 
paraffins.  The  most  fundamental  and  characteristic  aromatic 
hydrocarbon  is  benzene. 


42O  Introduction  to  General  Chemistry 


667.  Benzene,  CeEU. — -Benzene  is  a  colorless  liquid,  boiling 
at  79°  and  practically  insoluble  in  water.     It  is  a  by-product  of 
the  manufacture  of  coke  and  coal  gas  (634).    When  coal  is 
heated  in  the  absence  of  air  it  yields,  in  addition  to  coke  and 
gas,  a  large  amount  of  black  liquid  tar.    By  the  distillation  of 
tar  a  number  of  very  important  aromatic  hydrocarbons  are 
obtained.    One  of  the  most  useful  of  these  is  benzene.     We 
shall  give  the  structural  formula  of  benzene  without  attempting 
to  justify  our  reasons  therefor,  since  that  story  (although  a 
most  interesting  one)  is  entirely  too  long  for  this  text.     This 

formula  is 

H 

HC        CH 

CH 

\C/ 

H 

It  represents  a  ring  of  six  carbon  atoms,  with  three  single  and 
three  double  bonds,  and  six  hydrogen  atoms,  one  attached  to 
each  atom  of  carbon. 

Benzene  differs  from  the  paraffine  hydrocarbons  (which  are 
very  inactive)  in  being  remarkably  active  chemically.  This 
activity  is  shown  in  two  ways:  (i)  by  substitution  of  various 
radicals  for  one  or  more  hydrogen  atoms  of  each  molecule,  and 
(2)  by  addition  after  the  manner  of  unsaturated  compounds 
(660)  but  much  less  readily. 

668.  Other  Aromatic  Hydrocarbons. — Next  to  benzene  the 
simplest  aromatic  hydrocarbon  is  toluene,  or  methyl  benzene, 
C6H5CH3,  the  structural  formula  of  which  is 

H 

/c\ 

HC         CCH3 
HC         CH 

\c/ 

H 


Carbon  and  Carbon  Compounds  421 

In  order  to  save  labor  the  C6HS  radical  is  frequently  represented 
graphically  by  a  hexagon  called  the  benzene  ring,  so  that  the 
formula  of  toluene  is  written  thus: 

CH3 


Toluene  is  a  liquid,  boiling  at  110°  and  closely  resembling 
benzene.  It  is  obtained  from  coal  tar.  Xylene,  also  obtained 
from  the  same  source,  is  dimethyl  benzene,  C6H4(CH3)2.  In 
writing  the  structural  formula  of  xylene  we  note  that  there  are 
three  possible  arrangements: 

CH3  CH3 

CH3 


CH3  B 

CH3 

As  a  matter  of  fact  three  different  xylenes  are  known.  These 
are  called  ortho,  meta,  and  para  xylene,  respectively  written 
o-xylene,  m-xylene,  and  p-xylene. 

669.  Naphthalene,  CIOH8.  —  Naphthalene  is  a  white  crystalline 
substance  obtained  from  the  high-boiling  portion  of  coal  tar. 
It  is  extensively  used  in  the  household  under  the  name  of 
moth  balls.  Its  structure  is  represented  thus: 

H        H 


HC         C         CH 
CH 


HC         C 


H        H 

As  great  a  variety  of  organic  substances  are  derived  from 
naphthalene  as  from  benzene.  For  example,  we  have  two 
methyl  naphthalenes,  CIOH7  •  CH3: 

CH3 

k   CH3 


422  Introduction  to  General  Chemistry 

The  first  is  called  alpha  and  the  second  beta  methyl  naphthalene. 
Naphthalene  and  its  derivatives  are  important  starting  materials 
for  the  manufacture  of  dyestuffs. 

670.  Aromatic  Alcohols  and  Aldehydes. — Toluene,   CtH.s  • 
CH3,    can   be   considered  as  derived  from  methane,  CH4,  by 
the  substitution  of  the  phenyl  radical,  C6HS,  for  one  hydro- 
gen atom  of  CH4.     The  substitution  of  C6HS  for  H  in  methyl 
alcohol,  CH3  •  OH,  would  give  C6HS  •  CH2  •  OH.    The  sub- 
stance actually  exists  and  is  known  as  benzyl  alcohol.    Like 
other  alcohols,  it  can  be  oxidized  to  an  aldehyde  called  benzal- 

dehyde: 

C6HS  •  CH2  •  OH+O2->C6HS  •  CO  •  H+H2O. 

This  aldehyde  is  identical  with  the  principal  constituent  of 
the  oil  of  bitter  almonds.  Benzaldehyde  can  be  made  from 
toluene  by  converting  the  latter  into  a  chlorine  compound  and 
then  treating  this  product  with  water: 

C6H5  •  CH3+2C12->C6H5  •  CHCla+2HCl, 
C6H5  •  CHC12+H20->C6HS  •  CO  •  H+2HC1. 

Benzaldehyde  is  slowly  oxidized  by  contact  with  air  to  form 
benzoic  acid,  C6HS  •  CO  •  OH: 

2C6HS  •  CO  -  H+O2->2C6H5  •  CO  •  OH. 

671.  Benzoic  Acid,  C6H5  •  CO  •  OH.— This  important  acid 
is  a  white  crystalline  solid  which  can  be  made  in  several  ways  in 
addition  to  the  one  just  mentioned.     It  is  used  extensively  in 
the  form  of  its  sodium  salt,  sodium  benzoate,  C6H5CO  •  ONa, 
as  a  preservative  for  catsup  and  other  articles  of  food.     It  can 
be  used  legally  in  the  United  States  as  a  food  preservative  if  its 
presence  is  indicated  on  the  label.     It  occurs  naturally  as  a 
constituent  of  cranberries. 

Benzoic  acid  forms  salts  with  bases  of  all  kinds.  With 
alcohols  it  forms  esters.  The  latter  are  fragrant  liquids.  Ethyl 
benzoate,  CeHj  •  CO  •  OC2HS,  is  a  colorless  liquid,  boiling  at  211°. 

672.  Phenol,    C6HS  •  OH.— Phenol    (popularly    known    as 
carbolic  acid)  is  contained  in  coal  tar,  from  which  it  is  separated 
in  crude  form  by  distillation.     Phenol  is  also  made  on  a  large 


Carbon  and  Carbon  Compounds  423 

scale  from  benzene.  The  latter  substance  reacts  slowly  with 
sulfuric  acid,  forming  benzene  sulfonic  acid  and  water,  thus: 

C6H6+H2SO4  ->  C6HSSO2OH+H2O. 

The  product  is  an  acid  the  sodium  salt  of  which  when  fused  with 
sodium  hydroxide  gives  phenol : 

C6H5S02ONa+NaOH  ->  C6H5OH+Na2SO3. 

Phenol  is  a  white  crystalline  substance  having  a  peculiar,  char- 
acteristic odor.  It  is  moderately  soluble  in  water.  It  is  a 
violent  poison  and  is  extensively  used  as  a  germicide.  Phenol 
is  a  very  weak  acid  and  forms  salts  with  strong  bases : 

C6H5OH+NaOH  ->  C6H5ONa+H2O. 

It  is  interesting  to  contrast  phenol  and  ethyl  alcohol.     The 

formulae  of  the  two  substances 

H 

H    H  /C\ 

|       |  HC        COH 
H— C— C— OH 

II  HC        CH 
H    H  \c/ 

H 

show  that  both  contain  the  hydroxyl  radical  united  to  a  hydro- 
carbon radical.  In  consequence  we  might  expect  similar 
properties,  but  we  find  quite  the  contrary.  Phenol  shows  but 
few  of  the  characteristic  chemical  properties  of  an  alcohol. 

673.  Aromatic  Nitro  Compounds. — Aromatic  hydrocarbons, 
like  benzene,  react  very  readily  with  concentrated  nitric  acid 
in  a  peculiar  way,  as  illustrated  by  the  following  equation: 

C6H6+HN03  ->  C6HSNO2+H20. 

The  new  product  is  nitro  benzene,  a  light-yellow  colored  liquid 
of  aromatic  odor.  It  is  not  soluble  in  water  and  is  not  a  salt. 
The  structural  formula  is 


424  Introduction  to  General  Chemistry 

Toluene  and  nitric  acid  give  two  isomeric  nitro  compounds, 
CH3  •  C6H4  •  NO2.  These  have  the  following  formulae: 

CH3  CH3 

NOa 

NOa 

The  left-hand  formula  is  that  of  ortho  nitro  toluene ;  the  right- 
hand  one  that  of  para  nitro  toluene.  A  third  nitro  toluene  can 
be  made  by  indirect  methods.  This  substance,  called  meta 
nitro  toluene,  has  the  formula 

CH3 

NO3 

The  further  action  of  nitric  acid  on  either  ortho  or  para  nitro 
toluene  gives  finally  tri  nitro  toluene,  CH3  •  C6H2  -  (NO2)3. 
This  substance  is  the  violent  explosive  so  extensively  used  in 
the  war  and  known  popularly  as  T.N.T.  Its  formula  is 

CH3 

O.N/NNO, 

N02 

The  action  of  nitric  acid  on  phenol  gives  first  a  mixture  of 
ortho  and  para  nitro  phenol, 

OH 


Further  action  of  nitric  acid  finally  yields  tri  nitro  phenol,  or 
picric  acid,  C6H2(N03)3OH.  The  nitro  phenols  are  much 
stronger  acids  than  phenol  itself.  In  fact,  picric  acid  is  nearly 
as  strong  (highly  ionized)  an  acid  as  hydrochloric. 

Ammonium  picrate,  C6H2(NO3)3ONH4,  is  a  powerful  explo- 
sive.    It  has  been  extensively  used  in  the  war. 


Carbon  and  Carbon  Compounds  425 

674.  Aromatic  Amines. — Nitro  compounds  are  easily  acted 
on  by  reducing  agents,  as  illustrated  by  the  case  of  nitro  benzene: 

C6HS  •  NOa+3H2->C6Hs  •  NH2+2H20. 

The  new  product  is  called  aniline.  It  is  the  simplest  aromatic 
amine.  The  aromatic  amines  resemble  the  aliphatic  amines 
(658).  They  are  base-forming  substances  and  are  to  be  con- 
sidered as  ammonia  in  which  hydrogen  has  been  replaced  by  an 
aromatic  radical.  Aniline  unites  with  hydrochloric  acid  to 
form  a  true  salt,  a  chloride: 

C6HSNH2+HC1  ->  C6H5NH3C1. 

This  salt  corresponds  to  NH4C1.  A  great  variety  of  aromatic 
amines  can  be  made  by  the  reduction  of  nitro  compounds. 
Usually  the  reduction  is  carried  out  by  mixing  the  nitro  com- 
pound with  zinc  or  iron  and  hydrochloric  acid.  The  hydrogen 
liberated  by  the  action  of  the  metal  and  acid  then  reacts  with 
the  nitro  compound  in  the  way  above  indicated.  Aniline  and 
other  aromatic  amines  are  made  in  immense  quantities  to  be 
used  as  intermediates  in  the  manufacture  of  so-called  aniline 
or  coal-tar  dyes.  Further  reference  to  this  subject  will  be  found 
in  the  next  chapter. 


CHAPTER  XXVI 
ORGANIC  COMPOUNDS.    II 

675.  Introduction. — We  have  promised  to  show  the  reader 
some  of  the  successes  achieved  in  organic  chemistry  as  a  result 
of  the  systematic  study  of  the  science.     First  we  shall  take  up 
the  chemistry  of  foods  and  next  the  chemistry  of  explosives  and 
the  related  substances,  such  as  celluloid,  artificial  silk,  etc. 
After  this  we  shall  treat  briefly  the  poison  gases  used  in  war- 
fare, then  the  synthesis  of  essential  oils,  perfumes,  spices,  dyes, 
medicinals,  and  rubber.    This  may  seem  a  bewildering  list,  but 
organic  chemists  can  say  without  fear  of  contradiction  that 
they  have  accomplished  a  great  deal  more  than  is  even  casually 
mentioned  in  this  chapter.     Finally  we  shall  discuss  briefly  the 
chief  sources  of  materials  for  the  manufacture  of  organic  chem- 
icals. 

676.  The  Three  Classes  of  Foods. — All  foods  fall  into  three 
great  classes:    the  fats,  including  also  edible  oils;    the  carbo- 
hydrates, comprising  starches  and  sugars;  and  the  proteins,  in 
which  class  are  included  eggs,  lean  meat,  and  certain  nitrogenous 
vegetable  products.    A  well-balanced  diet  for  man  should  be 
made  up  of  foods  of  all  three  of  these  classes.    Fats  and  carbo- 
hydrates are  compounds  of  carbon,  hydrogen,  and  oxygen  only; 
while  all  proteins,  in  addition  to  these  three  elements,  contain 
also  nitrogen.    Of  the  three  classes  the  fats  are  from  the  chemical 
point  of  view  fhe  simplest,  and  their  chemistry  was  worked  out 
long  before  that  of  the  other  two  classes.    The  chemistry  of  the 
carbohydrates  was  well  cleared  up  during  the  last  two  decades 
of  the  nineteenth  century.    The  chemistry  of  the  proteins  is 
far  more  complex  and  is  even  today  far  from  completely  solved. 

677.  Fats. — The  term  fats  includes  liquid  as  well  as  solid 
animal  and  vegetable  products.    Liquid  fats  like  olive  oil,  cotton- 
seed oil,  peanut  oil,  castor  oil,  and  linseed  oil  are  chemically 
very  different  from  the  paraffin  or  mineral  oils  described  in  the 

426 


Organic  Compounds  427 

foregoing  chapter.  Most  natural  fats  (butter  fat  for  example) 
are  mixtures  of  several  chemical  compounds  all  of  which  belong 
to  a  single  group  of  organic  substances,  the  esters  (657).  The 
chemical  nature  of  these  esters  is  most  readily  shown  by  the 
conversion  of  fats  into  soap. 

678.  Soap. — Any  fat  is  changed  into  soap  when  it  is  boiled 
with  a  solution  of  sodium  hydroxide.     If  a  good  grade  of  white 
soap  is  dissolved  in  water  and  the  solution  acidified  with  hydro- 
chloric acid  a  dense  white  precipitate  forms.    The  evaporated 
filtrate  yields  only  common  salt.    The  white  precipitate  is  a 
mixture  of  three  or  four  fatty  acids  (655).    Among  the  com- 
monest fatty  acids  obtained  in  this  way  are  palmitic  acid, 
CISH3I  •  CO  •  OH,  and  stearic  acid,  CI7H3S  •  CO  •  OH. 

The  radicals  CI5H3I  and  CI7H3S  form  straight  or  unbranched 
carbon  chains,  as  illustrated  in  the  following  formula  for  palmitic 

acid: 

HHHHHHHHHHHHHHH 
HC-C-C'C-C-C'C-C-C-C'C-C-C-C-C-CO-OH 
HHHHHHHHHHHHHHH 

Stearic  acid  contains  one  more  CH2  group  per  molecule  than 
palmitic  acid. 

Soaps  are  the  sodium  (or  potassium)  salts  of  fatty  acids. 
The  action  of  HC1  on  sodium  palmitate  takes  place  thus : 

ClsH3ICOONa+HCl->CI5H3ICOOH+NaCl. 

679.  Glycerine    and    Its    Esters. — The    action    of    sodium 
hydroxide  on  a  fat  always  gives  in  addition  to  a  soap  one  other 
product,  glycerine.    Glycerine  is  a  sweet,  sirupy,  colorless  liquid, 
the  structural  formula  of  which  is 

CH2OH 
CHOH 
CH2OH 

Glycerine  is  an  alcohol,  but  it  differs  from  simple  (monatomic) 
alcohols  like  methyl  alcohol,  CH3OH,  and  ethyl  alcohol, 


428  Introduction  to  General  Chemistry 

CH3  •  CH2OH,  in  having  three  hydroxyl  groups  in  a  molecule. 
It  is  called  a  triatomic  alcohol.  It  will  be  recalled  (657)  that 
alcohols  are  not  bases,  in  spite  of  the  presence  of  hydroxyl 
groups.  They  do  not  yield  OH~  ions. 

Just  as  acetic  acid  and  ethyl  alcohol  unite  to  form  ethyl 
acetate  (an  ester)  and  water  (657),  so  a  fatty  acid  and  glycerine 
can  unite  to  form  an  ester  in  which  three  molecules  of  the  acid 
are  combined  with  one  of  glycerine.  Thus  palmitic  glycerine 
ester,  or  palmitin,  is 

CISH3ICOOCH2 

CISH3ICOOCH 
CISH3ICOOCH2 

This  substance  is  one  of  the  principal  constituents  of  beef  fat. 
Fats  in  general  are  the  glycerine  esters  of  various  fatty  acids. 
Just  as  ethyl  acetate  gives,  with  sodium  hydroxide,  sodium 
acetate  and  ethyl  alcohol, 

CH3  •  CO  •  OC2HS+ NaOH  -»  CH3  •  CO  •  ONa+ C2HSOH, 

so  a  fat  and  sodium  hydroxide  yield  a  soap  and  glycerine.  On 
account  of  the  close  chemical  relation  between  these  two  reactions 
the  first  as  well  as  the  second  is  spoken  of  as  a  saponification  of 
the  ester,  although  of  course  sodium  acetate  is  not  a  soap  in  the 
ordinary  sense  of  the  term. 

680.  The  Composition  of  Fats. — We  are  now  in  a  position  to 
understand  the  cause  of  the  differences  between  fats  from  various 
sources.  In  general,  a  given  sort  of  fat  is  a  mixture  of  the 
glycerine  esters  of  several  fatty  acids.  Among  such,  in  addition 
to  palmitic  and  stearic  acids,  already  mentioned,  we  have  oleic 
acid,  Ci7H33COOH;  lauric  acid,  CnH2ICOOH;  caprylic  acid, 
C7HISCOOH;  caproic  acid,  CsH^COOH;  valeric  acid,  C^COOH; 
and  butyric  acid,  C3H7COOH. 

Beef  fat  is  composed  largely  of  the  esters  of  palmitic,  stearic, 
and  oleic  acids.  These  esters  are  known  respectively  as  pal- 
mitin, stearin,  and  olein.  The  first  two  are  solids,  while  the 
last  is  an  oil  at  room  temperature.  Mutton  fat  resembles 


Organic  Compounds  429 

beef  fat  in  composition  but  contains  a  smaller  proportion  of 
olein,  while  hog  fat  in  the  form  of  lard  contains  appreciably 
more  olein  than  beef  fat.  Butter  fat  contains,  in  addition  to 
palmitin,  stearin,  and  olein,  a  considerable  proportion  of  butyrin, 
the  glycerine  ester  of  butyric  acid.  The  chief  constituent  of 
olive  oil  is  olein.  The  same  ester,  together  with  others,  com- 
prises cottonseed  oil,  an  edible  oil  of  enormous  economic 
importance. 

681.  The    Hardening   of    Oils.— Oleic   acid,    CI7H33COOH, 
differs  from  stearic  acid,  CI7H3SCOOH,  by  two  atoms  of  hydrogen 
per  molecule.    This  difference  is  the  result  of  one  double  bond 
(661,  662)  between  two  of  the  carbon  atoms  of   the  CI7H33 
radical,  which  is  therefore  an  unsaturated   (663)   compound. 
By  the  addition  of  hydrogen  to  the  double  bond,  oleic  acid  can 
be  converted  into  stearic  acid : 

CI7H33COOH+H2  ->  CI7H35COOH. 

By  a  similar  addition  of  hydrogen,  olein  is  changed  into 
stearin.  This  change  is  accomplished  by  the  aid  of  a  catalytic 
agent,  finely  divided  metallic  nickel.  By  means  of  this  process 
of  hydrogenation,  liquid  fats  like  cottonseed  oil  are  readily 
changed  into  solid  or  partially  solid  fats.  During  the  last  twenty 
years  this  so-called  hardening  of  fats  has  developed  into  an 
immense  industry.  The  product  made  from  cottonseed  oil  has 
about  the  consistency  of  lard  and  finds  extensive  use  as  a  sub- 
stitute for  the  latter. 

682.  The  Carbohydrates. — In  the  preceding  chapter  starch 
(636)  and  glucose,  or  grape  sugar  (639),  were  briefly  described. 
These  substances  belong  to  an  important  class  of  organic  com- 
pounds known  as  carbohydrates.     This  name  was  chosen  because 
these  substances  are  composed  of  carbon,  together  with  hydrogen 
and  oxygen,  the  two  latter  in  the  proportion  corresponding 
to  water.    Thus  glucose  is  C6HI2O6,  which  is  equivalent  to 
C6(H20)6.     However,  the  carbohydrates  are  not  simply  carbon 
with  water  of  hydra tion  in  the  sense  that  Na2S04  •  ioH20  is 
the  hydrate  of  Na2SO4. 


430  Introduction  to  General  Chemistry 

The  simplest  formula  which  would  represent  the  composition 
of  starch  is  C6HIOOS  (636);  but  it  is  certain  that  the  molecule 
of  starch  is  much  larger  than  that  represented  by  this  formula. 
The  formula  is  more  correctly  written  (C6HIOOS)W,  where  n  is 
an  integer  probably  as  large  as  30  or  40.  Ordinarily  the  simpler 
formula  is  employed. 

The  hydrolysis  of  starch  to  form  glucose  (638), 

(C6HI00S)M+ rcH20  ->  wC6HI206, 

is  a  very  important  reaction.  It  takes  place  readily  in  acid 
solution  by  reason  of  the  catalytic  action  of  H+  ions.  The 
higher  the  temperature  of  the  solution  the  more  rapid  the  hydra- 
tion  proceeds. 

Glucose  has  a  large  number  of  isomers  (650),  all  having,  of 
course,  the  same  formula.  These  sugars,  called  hexoses,  all 
have  properties  more  or  less  like  those  of  glucose.  Levulose, 
or  fruit  sugar,  is  one  of  the  commonest  of  the  hexoses.  It  is  the 
sugar  most  abundant  in  many  fruits. 

Ordinary  table  sugar,  commonly  known  as  cane  sugar  and 
called  by  chemists  sucrose,  has  the  formula  Ci2H22On.  It  is 
made  from  two  principal  sources,  sugar  cane  and  sugar  beets. 
It  has  the  same  composition  in  each  case.  Numerous  other 
plants  also  produce  sucrose.  Maple  sugar  is  largely  sucrose. 

Milk  sugar,  or  lactose,  C^H^On,  is  an  isomer  of  cane  sugar. 
It  is  present  in  cow's  milk  to  the  extent  of  4  per  cent.  It  is 
much  less  sweet  than  cane  sugar.  Maltose,  CI2H22On,  is  another 
isomer  of  cane  sugar.  It  is  formed  by  the  hydrolysis  of  starch 
in  the  presence  of  a  catalytic  agent  occurring  in  germinating 
seeds.  The  reaction  may  be  written 

2C6H10OS+H2O  ->  CI2H22OIZ, 
or  better, 

2  (C6HIOOS)W+  wH20  ->  wCI2H22Olx. 

The  catalytic  agent  is  called  diastase.  A  similar  substance, 
ptyalin,  is  present  in  saliva.  It  promotes  the  digestion  of  starch 
by  hydrolyzing  it  to  maltose.  The  catalytic  agents  diastase 
and  ptyalin  are  classed  as  enzymes. 


Organic  Compounds  431 

683.  The  Structure  of  the  Sugars. — The  structure  of  the 
simpler  sugars,  like  glucose,  was  worked  out  at  the  end  of  the 
nineteenth  century.  The  following  structure  of  glucose  was 
discovered  after  long  experimentation,  which  established  the 
presence  of  the  groups  indicated. 

H2COH 

HCOH 

I 
HCOH 

HCOH 
HCOH 
HCO 

The  five  hydroxyl  groups  behave  like  those  of  an  alcohol.  In 
this  respect  glucose  is  an  alcohol  somewhat  resembling  glycerine 
(679).  One  end  carbon  atom  of  the  glucose  molecule  forms  an 
aldehyde  radical  (652).  The  reactions  of  glucose  are  those  of 
an  alcohol  and  of  an  aldehyde.  Like  all  aldehydes  glucose  is 
a  good  reducing  agent.  It  reduces  an  alkaline  solution  of  copper 
to  cuprous  oxide.  This  reaction,  which  serves  as  the  best  test 
for  glucose,  is  carried  out  by  warming  .glucose  with  Fehling's 
solution.  This  solution  is  made  by  mixing  copper  sulfate  solu- 
tion with  a  solution  of  sodium  tartrate  (665)  containing  an 
excess  of  sodium  hydroxide.  In  the  presence  of  glucose  the 
deep-blue  Fehling's  solution  gives  a  red  precipitate  of  cuprous 
oxide,  Cu20. 

The  structure  of  levulose  is  represented  thus: 

H2COH 

HCOH 

HCOH 

HCOH 

CO 
H2COH 


432  Introduction  to  General  Chemistry 

It  is  an  alcohol  ketone  (656).  The  behavior  of  cane  sugar  with 
dilute  acids  throws  much  light  on  its  structure,  since  in  this 
reaction  it  unites  with  water  and  forms  equal  amounts  of  glucose 
and  levulose: 

CI2H22OXI+H2O  -»  CftHrA + C6HI206. 

From  this  it  follows  that  in  the  molecule  of  cane  sugar  a  molecule 
of  glucose  is  joined  with  one  of  levulose,  with  the  elimination  of 
a  molecule  of  water. 

Since  maltose  gives  by  hydrolysis  two  molecules  of  glucose 
its  molecule  may  be  considered  to  be  made  up  of  two  glucose 
radicals. 

Cane  sugar  and  maltose  do  not  reduce  Fehling's  solution. 
If  their  solutions  are  first  hydrolyzed  the  resulting  solution 
reduces  Fehling's  solution  readily. 

684.  Cellulose. — Cellulose,    which    occurs    nearly    pure    in 
cotton,  is  an  isomer  of  starch  (636).     Its  simplest  formula  is 
CeHjoOj,  but  its  true  formula   should  be  written  (C6H.I0Os)m, 
where  m  is  an  integer  probably  even  larger  than  n  in  the  starch 
formula.     Cellulose  is  far  less  active  chemically  than  starch  and 
is  practically  indigestible  by  man.     It  is  possible  to  hydrolyze 
cellulose  to  glucose,  but  the  reaction  takes  place  slowly.    Wood 
and  vegetable  fiber  in  general  contain  a  large  proportion  of 
cellulose.     The  latter  is  classed  with  starch  and  sugars  as  a 
carbohydrate. 

685.  The  Proteins. — We  shall  use  the  term  protein  to  include 
the  various  complex  nitrogenous  substances  forming  the  charac- 
teristic constituents  of  lean  meat,  white  of  eggs,  etc.    Most  plant 
seeds  also  contain  more  or  less  proteins.     Wheat  is  particularly 
rich  in  this  respect. 

Proteins  are  complex  substances  containing  in  addition  to 
carbon,  hydrogen,  and  oxygen  a  rather  large  percentage  of 
nitrogen  and  a  much  smaller  percentage  of  sulfur;  a  few  of  the 
proteins  also  contain  phosphorus.  Neither  the  exact  formula 
nor  the  structure  of  any  typical  protein  is  definitely  known.  The 
composition  of  albumin  (white  of  egg)  is  approximately  expressed 
by  the  formula  C720HII34N2I8S50248.  Although  considerable 


Organic  Compounds  433 

progress  has  been  made  in  recent  years  toward  the  elucidation 
of  the  structure  of  the  proteins,  much  remains  to  be  done.  When 
the  proteins  are  heated  with  acids  or  alkalies  they  are  split  up 
into  simpler  substances  which  are  found  to  be  nitrogen  deriva- 
tives of  fatty  acids  (655).  We  have  seen  that  acetic  acid, 
CH3COOH,  for  example,  forms  an  amide,  CH3CONH2,  acet- 
amide  (659).  We  have  also  learned  something  of  the  amines, 
of  which  methyl  amine,  CH3NH2,  is  the  simplest  representative, 
and  it  will  therefore  not  be  surprising  to  learn  that  a  substance 
having  the  formula 

H2NCH2COOH, 

which  we  may  call  amino  acetic  acid,  or  glycocoll,  exists.  This 
acid  can  form  an  amide, 

H2NCH2CONH2, 

which  can  unite  with  one  or  more  molecules  of  glycocoll  to  form 
such  products  as 

H2NCH2CONHCH2CONH2 
and 

H2NCH2CONHCH2CONHCH2CONH3. 

Still  more  complex  substances,  amino  acids,  have  been  built 
up  artificially  in  the  laboratory.  The  fact  that  these  substances 
are  identical  with,  or  related  to,  the  decomposition  products  of 
the  proteins  leads  us  to  think  that  the  molecules  of  the  latter 
are  made  up  of  amino  acid  radicals,  among  which  are  those  of 
fatty  acids  other  than  acetic.  Proteins  from  different  sources 
differ  markedly  from  one  another  by  reason  of  the  kinds  and 
relative  amounts  of  the  amino  acids  which  compose  them. 

Many  proteins  are  of  vegetable  origin.  The  cereals  like 
wheat,  oats,  rye,  barley,  and  corn  are  comparatively  rich  in 
these  nitrogenous  substances  (511).  Beans,  peas,  and  other 
legumes  also  contain  large  percentages  of  proteins,  while  vege- 
tables contain  but  small  amounts. 

686.  Why  the  Body  Needs  Food.— The  body  needs  food  for 
growth,  repair,  and  the  supply  of  energy.  For  the  adult  only 


434  Introduction  to  General  Chemistry 

the  last  two  are  of  importance.  The  average  amount  of  carbon 
dioxide  exhaled  per  day  by  an  adult  is  about  1,100  grams;  but 
the  amount  varies,  greatly  increasing  with  the  amount  of  work 
done.  Nitrogen  is  excreted  largely  as  urea  (659)  but  also  in 
smaller  amounts  in  the  form  of  other  compounds.  The  average 
daily  loss  calculated  as  nitrogen  amounts  to  about  20  grams  for 
an  adult.  Hydrogen  and  oxygen  are  also  lost  in  large  quantities, 
principally  in  the  form  of  water  and  to  a  smaller  extent  in  com- 
pounds with  carbon,  nitrogen,  etc.  These  losses  must  be  com- 
pensated by  food  and  water  in  order  to  maintain  bodily  weight; 
for  of  course  the  law  of  the  conservation  of  matter  (21)  applies 
rigorously  to  all  bodily  processes. 

687.  The  Law  of  the  Conservation  of  Energy  for  Bodily 
Processes. — The  body  expends  energy  in  two  ways:  in  doing 
work  and  in  giving  off  heat  to  the  surroundings.  The  source  of 
this  energy  is  found  in  the  chemical  changes  of  the  food  eaten 
and  the  oxygen  inhaled.  The  amount  of  energy  produced  when 
a  known  amount  of  a  given  foodstuff,  together  with  sufficient 
oxygen,  is  changed  to  the  same  products  as  those  formed  in  the 
body  'can  be  determined  by  means  similar  to  that  described 
earlier  (357).  The  energy  per  gram  of  a  food  may  be  expressed 
in  calories  and  called  its  fuel  value. 

Until  recent  years  it  was  a  question  whether  the  amount  of 
energy  supplied  by  the  food  eaten  was  exactly  equal  to  that 
expended  in  the  form  of  work  plus  that  given  off  as  heat  when 
the  body  neither  gained  nor  lost  in  weight.  This  important 
problem  was  solved  by  the  very  elaborate  experiments  of  the 
American  chemists  Atwater  and  Benedict.  These  scientists 
constructed  a  huge  calorimeter  (357)  in  which  a  man  could  live 
and  perform  work  for  hours  at  a  time.  The  amount  of  work 
done  and  heat  given  off  was  accurately  measured,  and  the 
energy  or  calorific  value  of  all  food  eaten  was  determined. 
After  several  years  of  the  most  careful  work,  in  which  animals 
as  well  as  men  were  experimented  upon,  it  was  conclusively 
proved  that  the  energy  given  out  by  the  body  is  exactly  equal 
to  that  produced  when  the  food  and  oxygen  taken  are  changed 
to  the  same  forms  as  those  excreted  by  the  body.  In  other 


Organic  Compounds  435 

words,  it  was  found  that  the  law  of  the  conservation  of  energy 
applies  rigidly  to  all  bodily  processes. 

688.  The  Science  of  Dietetics.— The  facts  set  forth  in  the 
two  preceding  sections  serve  as  the  basis  for  a  scientific  treat- 
ment of  the  subject  of  nutrition.    This  branch  of  science  is 
called  dietetics.      A  satisfactory,   well-balanced  ration  must 
supply  sufficient  amounts  of  each  class  of  food  to  compensate 
for  the  known  or  determinable  body  losses  of  material  and  at 
the  same  time  yield  sufficient  energy  to  enable  the  performance 
of  the  required  amount  of  physical  work  and  also  keep  up  bodily 
temperature,  all  without  the  loss  of  body  weight. 

The  energy  requirement  of  a  man  doing  moderate  work  is 
about  3,200  kilogram  calories.1  To  get  this  he  must  consume 
food  more  than  sufficient  to  satisfy  his  requirements  for  carbon; 
but  this  will  not  necessarily  also  supply  his  requirements  for 
nitrogen.  Therefore  a  sufficient  ration  will  result  from  eating 
enough  protein  to  compensate  the  nitrogen  loss,  and  in  addition 
enough  fats  or  carbohydrates,  or  better  both,  to  bring  the  fuel 
value  up  to  his  requirement.  The  ratio  of  carbohydrates  to 
fat  in  the  diet  is  not  of  fundamental  importance;  furthermore 
the  protein  consumption  may,  for  some  individuals  or  even  for 
whole  races,  be  safely  increased  far  above  the  necessary  minimum. 
For  adult  male  Americans,  professional,  business  men,  and 
students,  loog.  of  protein,  i25g.  of  fat,  and  400  g.  of  carbo- 
hydrates, with  a  total  fuel  value  of  3,200  kilo  calories,  consti- 
tutes an  average  daily  ration. 

Table  XXXII  gives  the  data  for  sample  meals  for  one  day 
for  one  man  doing  moderately  heavy  work.  It  shows  the 
weight  in  fractions  of  a  pound  and  in  grams  of  each  article,  and 
also  its  protein  content  and  fuel  value. 

689.  Vitamines. — It  would  naturally  be  inferred  from  the 
discussion  of  the  preceding  sections  that  in  providing  for  a 
dietary  sufficient  to  maintain  health  and  weight  it  is  necessary 
to  consider  only  the  protein  content  and  fuel  value  of  the  food 

'On  account  of  the  fact  that  the  calorie  (ux)  is  so  small  a  unit  of  heat,  a 
unit  one  thousand  fold  larger  is  in  common  use.  This  larger  unit  is  the  heat 
required  to  raise  the  temperature  of  one  kilogram  of  water  one  degree.  It  may  be 
called  a  kilogram  calorie,  or  kilo  calorie. 


436 


Introduction  to  General  Chemistry 


supplied.     Yet  it  has  long  been  known  that  in  earlier  times 
sailors  who  lived  for  long  periods  on  an  abundance  of  food  of 


TABLE  XXXII 


ARTICLE 

POUND 

GRAMS 

PERCENT- 
AGE OF 
PROTEIN 

PROTEIN 
IN  GRAMS 

FUEL  VALUE 
IN  KILO 
CALORIES 

Breakfast 

Bread         

o  20 

no 

0    2 

8   28 

2AO 

Half  an  orange  
Sugar 

0.30 
One 

135 

22 

0.6 
O    O 

0.81 

O    OO 

50 

88 

Two  eggs  

O    2O 

QO 

ITT 

II    70 

1  60 

Butter 

O   O? 

23 

I    O 

O    21 

I7O 

Cup  of  coffee  and  cream  .  .  . 

O.O5 

90 

Totals  

21    l6 

708 

Luncheon 

Tomato  soup  

O    2O 

OO 

j    i 

I    OO 

sO 

Lamb  chops  

O    2O 

OO 

I  •?    r 

12    22 

281 

Peas  

O    12 

c.6 

8   0 

4ro 

66 

Stewed  apples  
Sugar                        

O.  12 

56 

0-3 

0.17 

55 
88 

Milk                    

^6 

Bread  
Butter  

O.  2O 

go 

•6 
9.2 
I    O 

8.28 
O    21 

A55 

240 

u.  13 

Totals  

7.2     80 

I   IO7 

Dinner 

Consomme 

O.  2O 

go 

Roast  beef  

0.40 

181 

I5-° 

27.00 

400 

Baked  potato  
Cauliflower  
Lettuce 

0.50 
0.12 
O    12 

227 
56 
56 

1.8 

i-4 

I    O 

4.09 
0.78 

o  60 

197 
15 

IO 

French  dressing  

O.O3 

16 

IOO 

Bread  
Butter  
Ice-cream  

O.  2O 
0.05 
O.  2O 

90 

23 
90 

9.2 

i  .0 

2.6 

8.28 
0.23 

2.40 

240 

170 

1  60 

Black  coffee 

O    IO 

AC 

Total  

43  •  38 

1,292 

Total  for  the  day 

1     08.43 

7,107 

1 

! 

very  limited  variety,  such  as  salt  pork  and  "hard- tack,"  were 
liable  to  be  afflicted  with  a  peculiar  and  often  fatal  disease  called 
scurvy.  Afflicted  persons  rapidly  recovered  when  supplied  with 


Organic  Compounds 


437 


fresh  vegetables  or  even  with  the  juice  of  oranges,  lemons,  or 
limes.  The  British  navy  and  mercantile  marine  have  for  fifty 
years  or  more  been  required  to  provide  sailors  regularly  with 
lime  juice.  Scurvy  is  now  of  rare  occurrence. 

A  still  more  remarkable  case  is  found  in  the  cause  of  and 
remedy  for  the  peculiar  oriental  disease  beri-beri.  This  affects 
people  who  live  chiefly  on  a  diet  of  rice  and  fish.  Recent 
experiments,  particularly  with  pigeons,  has  shown  that  these 
birds  thrive  on  natural  rice,  while  if  fed  on  the  polished  grains 
they  quickly  suffer  from  the  disease  and  soon  die.  Dangerously 
sick  pigeons  make  marvelous  recovery  when  given  small  amounts 
of  the  material  removed  from  the  grains  in  the  process  of  polish- 
ing. Further  experimentation  has  shown  conclusively  that 
beri-beri  is  the  result  of  a  diet  deficient  in  a  substance  known 
as  water  soluble  B  found  in  the  germ  of  the  rice  grain.  This 
same  substance  is  also  found  in  the  germ  of  wheat  and  other 
grains. 

Very  extensive  experiments  have  shown  conclusively  that 
rats  cannot  grow  nor  even  live  long  on  diets  containing  adequate 
amounts  of  proteins,  fats,  and  carbo- 
hydrates if  every  article  of  food  has 
been  highly  purified  by  chemical  pro- 
cesses. This  is  because  certain  essential 
substances  contained  in  the  natural 
foods  have  been  removed  in  the  pro- 
cesses of  purification.  Milk  is  espe- 
cially rich  in  these  essential  substances. 
By  adding  small  amounts  of  milk  to 
the  diet  of  rats  living  on  purified  food  a 
wholly  satisfactory  ration  is  obtained. 
The  results  of  such  experiments  are 
shown  in  Fig.  97.  The  lower  curve 
shows  the  change  in  average  weight  of 
six  young  rats  fed  on  a  diet  of  purified  foods  alone.  The  upper 
curve  shows  the  weights  of  the  same  number  of  young  rats 
which  received  the  same  food  ration  as  the  first  six  but  had  in 
addition  2  c.c.  each  of  milk  per  day.  Work  of  F.  G.  Hopkins. 


o 


\ 


20 

Days 


FIG.  97 


438  Introduction  to  General  Chemistry 

In  addition  to  water  soluble  B,  milk  contains  also  another 
peculiar  substance  called  fat  soluble  A.  The  latter  occurs  also 
in  the  leafy  parts  of  most  vegetables.  The  term  vitamines 
is  usually  used  to  designate  the  important  substances,  fat 
soluble  A  and  water  soluble  B.  It  is  now  definitely  established 
that  animals,  including  man  (especially  children),  must  have 
for  growth  and  health  a  constant  supply  of  both  these  vitamines. 
A  diet  (for  man)  of  roots,  tubers,  seeds,  and  meat  may  furnish 
sufficient  proteins  and  have  adequate  fuel  value  and  still  be 
markedly  deficient  in  vitamines,  particularly  fat  soluble  A. 
The  deficiency  is  best  avoided  by  the  liberal  use  of  milk  and  leafy 
vegetables;  of  these,  milk  is  the  more  important. 

The  chemical  nature  of  the  vitamines  still  remains  to  be 
discovered,  but  their  importance  as  food  constituents  is  no  longer 
in  question.  There  is  little  doubt  that  vitamines  act  cata- 
lytically. 

690.  Dependence   of  Animals  upon  Plants. — All  animals, 
including  man,  depend  for  their  food  upon  plants  or  upon  other 
animals  which  in  turn  feed  upon  plants.    Animals  cannot  sub- 
sist upon  the  free  elements  that  compose  their  food,  nor  even 
upon  the  simpler  compounds  of  these  elements,  such  as  carbon 
dioxide,  hydrocarbons  (643),  ammonia,  amines  (658),  etc.,  but 
must  have  the  far  more  elaborate  compounds,  the  carbohydrates 
(682),  fats  (677),  and  proteins  (685),  and  in  addition  mineral 
salts  and  vitamines  (689). 

Plants,  on  the  other  hand,  require  for  their  sustenance  far 
simpler  materials,  principally  carbon  dioxide,  water,  and  simple 
nitrogen  compounds,  such  as  ammonia  or  nitrates,  and  also 
small  amounts  of  mineral  matter.  Some  of  the  simplest  organ- 
isms contain  species  of  bacteria,  by  aid  of  which  they  assimilate 
free  nitrogen  from  the  air  (515). 

691.  Photosynthesis. — The  plant  products  formed  from  these 
simple  inorganic  substances  have  far  more  energy  than  the 
latter.    What  then  is  the  source  of  this  energy?    Plainly  the 
light  and  heat  necessary  for  the  growth  of  all  plants  (excepting 
fungi  and  other  parasitic  plants  which  live  on  decaying  animal 
or  vegetable  matter).    The  energy  of  sunlight  is  transformed 


Organic  Compounds  439 

in  the  growing  plant  into  the  chernical  energy  of  the  plant 
products. 

Cellulose  and  starch  are  the  most  abundant  plant  products. 
For  the  formation  of  these  only  carbon  dioxide  and  water  are 
theoretically  required : 

6COa+  sH20  ->  C6HI0Os-{-6O2. 

However,  this  reaction  does  not  take  place  in  the  simple  fashion 
indicated  by  this  equation.  It  takes  place  only  in  plants  and 
then  only  in  such  parts  as  contain  the  characteristic  green  sub- 
stance chlorophyl.  Probably  several  intermediate  stages  exist 
in  the  change  of  carbon  dioxide  into  starch  or  cellulose;  in  any 
case  oxygen  is  always  a  product  of  plant  growth.  Since  this 
building  up  of  complex  products  occurs  by  the  aid  of  light  the 
process  is  termed  photosynthesis. 

692.  Some  Common  Organic  Explosives. — We  have  referred 
earlier  (571)  to  the  great  practical  importance  of  explosives. 
The  most  useful  explosives  are  made  by  the  nitration  of  organic 
substances  such  as  glycerine  (679),  cotton  (684),  phenol  (672), 
and  toluene  (668).  The  products  thus  obtained  will  now  be 
briefly  described. 

Nitroglycerine  is  formed  when  glycerine  is  added  drop  by 
drop  to  a  cooled  mixture  of  concentrated  nitric  and  sulfuric 
acids.  The  nitroglycerine  separates  as  an  insoluble  heavy  oil 
when  the  product  is  poured  into  water.  It  is  a  violent  explosive 
which  finds  extensive  use  in  blasting.  Nitroglycerine  is  the  nitric 
acid  ester  (657)  of  the  alcohol  glycerine  (679).  Its  formula  is 

H2CONO2 
HCONO2 
H2CONO2 

Its  explosive  nature  arises  from  the  fact  that  it  contains  more 
than  sufficient  oxygen  to  change  all  its  hydrogen  and  carbon  into 
water  and  carbon  dioxide  respectively.  At  the  high  tempera- 
ture reached  in  the  explosion  water  is,  of  course,  gaseous,  as  are 
likewise  the  other  products,  CO2  and  N2.  Therefore  the  products 


440  Introduction  to  General  Chemistry 

of  the  explosion  occupy  many  times  the  volume  of  the  original 
substance.  This  explains  the  enormous  force  produced  by  the 
explosion.  Dynamite,  which  is  a  solid  mixture  of  infusorial 
earth  and  nitroglycerine,  is  much  safer  to  handle  than  the  latter 
substance. 

693.  Nitrocellulose,  or  guncotton,  is  made  from  cotton  by 
a  process  similar  to  that  used  in  making  nitroglycerine.  The 
product  is  a  white  solid  scarcely  differing  in  appearance  from  the 
cotton  from  which  it  is  made.  It  is  far  less  sensitive  to  shock 
than  nitroglycerine,  especially  when  in  a  moist  state,  and  since 
it  can  be  perfectly  exploded  while  moist  by  the  use  of  a  detonator 
(573)  it  is  one  of  the  safest  and  most  useful  of  explosives.  Gun- 
cotton  can  be  physically  compounded  with  nitroglycerine  to 
form  cordite,  a  transparent  solid  resembling  amber  in  appear- 
ance. This  is  one  of  the  most  important  propellants  for  large 
projectiles.  Modern  smokeless  powders  are  products  closely 
related  to  guncotton. 

The  complete  nitration  of  phenol,  C6HSOH  (672),  yields 
tri  nitro  phenol,  or  picric  acid,  C6H2(NO3)3OH.  This  substance 
and  its  ammonium  salt  (673)  are  powerful  explosives.  Their 
most  extensive  use  is  in  shrapnel. 

The  explosive  commonly  designated  as  T.N.T.  is  tri  nitro 
toluene.  It  has  also  been  mentioned  earlier  (571).  During 
the  war  it  was  used  in  enormous  amounts  in  shrapnel  and  other 
explosive  shells. 

A  few  words  may  be  added  at  this  point  regarding  the  differ- 
ences in  explosives,  since  the  student  will  naturally  wonder  why 
so  many  different  explosives  are  used.  The  most  important 
properties  which  determine  the  character  of  an  explosive  are 
(i)  its  sensitiveness,  (2)  its  force  of  explosion,  determined  by 
the  volume  and  temperature  of  the  gaseous  products,  and 
(3)  its  velocity  of  explosion.  Of  these  three  properties  the  last 
is  very  important;  for  although  ordinary  observation  would 
indicate  that  every  explosion  is  instantaneous,  this  is  far  from 
being  the  case.  Every  explosion  requires  time  for  its  comple- 
tion. A  satisfactory  propellant  must  not  explode  too  rapidly. 
It  must  allow  time  for  the  projectile  to  get  under  way,  otherwise 


Organic  Compounds  441 

it  would  burst  the  gun.  For  exploding  shrapnel  shells  and  for 
blasting  rocks  rapidly  exploding  substances  are  used. 

694.  Other  Products  of  Nitrocellulose. — When  the  nitration 
of  cotton  is  not  carried  so  far  as  in  the  preparation  of  guncotton 
the  product  is  known  as  pyroxyline  or  soluble  cotton.  This 
substance  resembles  guncotton  closely  but  is  less  explosive. 
It  dissolves  readily  in  many  organic  solvents,  such  as  acetone 
(656)  and  esters  (657),  and  in  a  mixture  of  alcohol  (641)  and 
ether  (642).  The  solution  so  obtained  is  called  collodion;  it  is 
used  as  an  adhesive,  as  a  liquid  court-plaster,  as  a  coating  for 
incandescent  gas  mantles,  and  for  many  other  purposes. 

Artificial  silk,  a  product  resembling  silk  in  appearance,  but 
totally  different  chemically,  is  made  in  one  way  from  collodion. 
The  process  consists  in  "spinning"  a  concentrated  collodion 
solution  from  a  fine  glass  capillary  by  means  of  high  pressure. 
Upon  coming  into  the  air  the  solvents  evaporate  at  once,  leaving 
a  filament  of  the  fineness  of  a  natural  silk  fiber.  These  filaments 
are  made  into  threads ;  but  this  material  is  extremely  combustible 
and  must  be  denitrated  by  a  chemical  process  which  changes 
it  back  into  cellulose  without  altering  its  beautiful  silky  luster. 
Artificial  silk  is  made  in  large  quantities. 

A  rough  way  of  telling  the  difference  between  artificial  and 
natural  silk  is  to  set  fire  to  a  small  piece  of  the  fabric.  Natural 
silk  burns  the  way. hair  does,  melting  back  of  the  flame  and 
giving  off  a  characteristic  odor.  Artificial  silk  burns  as  does  a 
thin  piece  of  cotton.  It  does  not  melt  and  gives  virtually  no  odor. 

Celluloid  is  made  by  thoroughly  kneading  and  rolling  a  warm 
mixture  of  pyroxyline  and  camphor,  CIOHl60.  The  latter  is  a 
product  of  the  camphor  tree.  The  many  forms  and  uses  of 
celluloid  are  too  well  known  to  require  description.  Photo- 
graphic films  made  from  this  material  are  highly  inflammable. 
On  this  account  celluloid  is  combined  with  less  inflammable 
materials  like  cellulose  acetate  (acetic  acid  ester  of  cellulose)  in 
the  preparation  of  photographic  films  for  moving  pictures. 

Leather  substitute,  or  artificial  leather,  is  made  by  coating 
heavy  cotton  cloth  with  a  preparation  in  which  pyroxyline  is 
the  principal  ingredient. 


442  Introduction  to  General  Chemistry 

695.  The  Materials  of  Chemical  Warfare. — The  introduction 
by  the  Germans  of  chlorine  in  warfare  was  rapidly  followed  by 
the  use  of  several  other  poisonous  and  irritating  substances, - 
practically  all  of  which  were  carbon  compounds.  One  of  the 
most  important  of  these  was  phosgene,  or  carbon  oxychloride, 
COC12.  This  gas  is  made  by  the  union  of  carbon  monoxide 

(632)  and  chlorine: 

CO+C12-»COC12 

It  has  a  powerful,  choking  odor  and  is  very  poisonous.  It  is 
readily  liquefied  (boiling-point  8°).  It  was  frequently  mixed 
with  liquid  chlorine  in  gas  attacks. 

The  so-called  mustard  gas  is  not  a  gas  at  all  but  a  colorless 
liquid  known  to  chemists  as  dichlor  diethyl  sulfide  (C1C2H4)2S. 
It  has  a  rather  faint  odor  and  at  first  appears  harmless  enough; 
but  inhalation  of  the  vapor  or  mist  produced  by  explosion  of  a 
shell  containing  some  of  it  causes  terrific  inflammation  of  the 
lungs,  frequently  proving  fatal.  Almost  inconceivably  minute 
amounts  on  the  skin  cause  in  the  course  of  a  few  days  deep  and 
dangerous  wounds.  The  substance  is  made  from  ethylene,  C2H4, 
and  sulfur  chloride,  S2C12  (601),  and  has  the  following  structure: 

H    H          H    H 

Cl— C— C— S— C— C— Cl 

I  j 

H    H          H    H 

It  seems  to  act  by  being  absorbed  by  the  skin  and  then  slowly 
hydrolyzing  within  the  tissues  to  form  hydrochloric  acid  and 
other  products. 

Chloropicrin,  CC13N02,  made  by  the  action  of  bleaching 
powder  (351)  on  picric  acid  (673),  is  a  high-boiling  liquid.  It 
is  extremely  irritating  to  the  eyes,  is  poisonous,  and  causes 
vomiting.  It  passes  through  clothing  and  the  fabric  of  a  gas 
mask  rather  readily,  and  it  is  difficultly  absorbed  in  the  canister 
of  chemicals  used  with  a  mask.  This  substance  was  used  in 
large  amounts  toward  the  end  of  the  war.  Chloropicrin  had 
no  use  before  the  war;  it  is  now  proposed  to  employ  it  as  an 
insecticide. 


Organic  Compounds  443 

696.  Organic  Synthesis. — The  term  organic  synthesis  means 
the  artificial  building  up  of  organic  compounds  from  simpler 
substances,  ultimately  from  the  elements.     A  century  ago  it  was 
generally  believed  that  the  complex  carbon-containing  products 
of  plants  and  animal  organisms   (organic   compounds)   were 
formed  as  the  result  of  vital  forces,  and  that  it  was  impossible 
for  chemists  to  make  them  artificially.     In  1828  Wohler  proved 
the  fallacy  of  this  idea  by  the  synthesis  of  one  of  the  most  char- 
acteristic   of   animal   products,    urea,    CO(NH2)2    (659).     His 
method  consisted  in  oxidizing  potassium  cyanide,  KNC  (665), 
to  the  cyanate  KNCO  by  heating  it  with  litharge,  PbO: 

KNC+PbO->KNCO+Pb. 

From  KNCO  and  (NH4)2SO4  he  obtained  by  double  decomposi- 
tion ammonium  cyanate,  NH4NCO.  This  salt  when  -warmed 
in  solution  gradually  changes  into  urea, 

NH4NCO->CO(NH2)2. 

The  announcement  of  Wohler's  discovery  created  a  profound 
sensation  and  led  to  the  expectation  that  other  vital  products 
could  also  be  synthesized.  This  expectation  has  in  the  past 
ninety  years  been  realized  far  more  completely  than  the  most 
enthusiastic  chemist  of  Wohler's  day  could  have  predicted;  for 
not  only  have  the  greatest  variety  of  vital  products  been  synthe- 
sized, but  thousands  of  related  and  new  and  unrelated  organic 
substances  have  been  made  in  the  laboratory,  so  that  today  their 
number  is  legion.  True,  however,  much  remains  to  be  done, 
for  as  we  have  seen  in  one  instance  the  proteins  have  as  yet 
not  been  synthesized.  By  way  of  illustrating  the  nature  and 
scope  of  organic  synthesis  we  shall  in  the  following  paragraphs 
recount  briefly  a  few.  typical  cases. 

697.  Essential    Oils    and    Perfumes. — Spices,    fruits,    and 
flowers  owe  their  characteristic  odors  or  perfumes  to  small 
amounts  of  substances  which  can  in  many  cases  be  isolated  and 
purified.     These  fragrant  substances  are  volatile  oils  or  solids 
known  generally  as  essential  oils.    Frequently  the  oil  from  a 
given  source  is  a  mixture  of  several  definite  chemical  substances, 
although  in  other  cases  but  a  single  substance  is  present. 


444  Introduction  to  General  Chemistry 

The  first  essential  oil  made  synthetically  was  that  contained 
in  bitter  almonds  and  well  known  as  the  oil  of  bitter  almonds. 
Examination  of  the  natural  oil  showed  it  to  be  benzaldehyde,  a 
substance  having  the  formula 

H 

:o 


When  toluene  (668)  is  treated  with  chlorine  it  gives  benzal 
chloride,  C6HSCHC12.  The  latter,  by  treatment  with  water, 
gives  benzaldehyde  (670).  The  product  so  obtained  has  the 
same  agreeable  odor  as  that  made  from  bitter  almonds  and  is 
extensively  used  as  a  substitute  for  the  natural  oil  in  flavoring 
extracts. 

The  chief  constituent  of  oil  of  cinnamon,  known  as  cinnamic 
aldehyde,  has  the  formula 

C6H5CH  =  CHCHO. 

It  can  also  be  made  synthetically  from  toluene,  and  when  so 
prepared  has  the  same  fragrant  odor  as  cinnamon. 

The  characteristic  constituent  of  vanilla  is  also  an  aldehyde, 
called  vanillin,  somewhat  related  to  the  two  foregoing  substances. 

Its  formula  is 

H 


It  is  a  white  crystalline  solid,  which  is  present  in  vanilla  beans 
to  the  extent  of  about  i  per  cent.  After  the  discovery  of  its 
formula  as  just  given  it  became  possible  to  make  the  substance 
synthetically.  It  was  soon  found  that  other  essential  oils  were 
closely  related  to  vanillin.  Thus  oil  of  cloves  consists  largely  of 


H2CH  = 


Organic  Compounds  445 

This  oil  is  easily  and  cheaply  made  from  cloves.  It  can  be 
converted  into  vanillin  by  a  process  of  oxidation,  which  thus 
affords  a  good  practical  method  of  making  the  latter  valuable 
substance. 

It  is  a  well-known  fact  that  organic  substances  with  closely 
related  formulae  have  similar  properties,  and  this  is  beautifully 
illustrated  by  the  near  relatives  of  the  two  substances  just  con- 
sidered. The  fragrant  oil  of  sassafras  consists  largely  of 
safrole, 

_'0/\CHaCH=CHa 

H2 


This  substance  can  be  oxidized  to  an  aldehyde  piperonal, 

/VH 

HC/°I        |C° 

H<c\olJ 

a  substance  having  the  delightful  odor  of  heliotrope.  Thus  an 
exquisite  and  costly  perfume  is  made  in  the  laboratory  from  the 
cheap  and  abundant  oil  of  sassafras. 

Many  essential  oils  belong  to  the  class  of  substances  known 
as  esters  (657).  For  example,  oil  of  wintergreen  is  the  methyl 
ester  of  salicylic  acid,  readily  made  by  following  reaction: 

/\OH 
+  CH3OH^|  „  +H,0 


*r*r\r\r* 
v     /LU(JC 


H 


The  ester,  which  is  made  technically  on  a  large  scale,  is  a  color- 
less oil  having  exactly  the  odor  of  wintergreen.  It  will  be  of 
interest  to  trace  the  manufacture  of  this  substance  from  its 
beginning.  To  get  salicylic  acid  we  start  with  benzene  (667), 
obtained  from  coal  tar.  This  by  treatment  with  sulfuric  acid 
gives  benzene  sulfonic  acid  (672), 


/\S02OH 


446  Introduction  to  General  Chemistry 

When  the  sodium  salt  of  this  acid  is  fused   with  an  excess  of 
sodium  hydroxide  it  yields  phenol  or  carbolic  acid  (672), 


The  sodium  salt  of  this  substance  reacts  with  carbon  dioxide  to 
form  the  salt  of  salicylic  acid, 

C6H5ONa+  CO2  ->  C6H4OHCOONa, 

from  which  the  free  acid  is  readily  obtained  by  the  action  of 
sulfuric  acid. 

The  subjects  discussed  in  this  section  are  of  importance 
chiefly  as  a  means  of  illustrating  the  ways  in  which  some  natural 
products  are  made  in  the  laboratory.  It  is  not  worth  while  for 
the  student  to  attempt  to  memorize  the  formulae  or  the  reactions 
involved.  The  examples  given  are  chosen  from  hundreds 
equally  important  and  interesting. 

698.  Dyes.  The  Synthesis  of  Indigo. — In  earlier  times 
dyes  were  usually  plant  or  animal  products.  A  few  coloring 
substances,  not  dyes  in  the  strict  sense  of  the  word,  were  of 
mineral  origin.  At  present  all  but  a  small  proportion  of  dyes 
are  made  artificially  from  substances  obtained  from  coal  tar. 
In  a  few  cases  dyes  originally  obtained  from  plants  are  now 
made  in  the  laboratory.  The  most  interesting  case  in  point  is 
that  of  indigo,  which  has  long  been  one  of  the  most  important 
blue  dyes.  It  is  the  product  of  a  plant  grown  extensively  in 
India,  Java,  and  elsewhere.  After  the  development  of  synthetic 
chemistry  it  was  apparent  that  the  artificial  production  of  any 
natural  chemical  substance  was  possible,  though  doubtless  in 
many  cases  difficult  of  realization.  This  possibility  led  chemists 
to  hope  that  indigo  could  be  made  in  the  laboratory.  The  first 
step  was  to  discover  its  structural  formula.  The  purification 
of  indigo  proved  easy,  and  the  analysis  of  the  pure  crystalline 
substance  was  a  matter  of  routine.  The  simplest  formula 
possible  according  to  the  analysis  was  C8HSNO;  but  the  vapor 
density  (71,  217)  indicated  a  molecular  weight  corresponding 
to  twice  this  formula,  namely,  Cl6HION2O2.  Now  it  would  be 


Organic  Compounds  447 

possible  to  think  of  hundreds  of  molecular  arrangements  of 
1 6  atoms  of  carbon,  10  atoms  of  hydrogen,  and  2  each  of  nitrogen 
and  oxygen,  all  in  accord  with  accepted  laws  of  valence  (648) ; 
only  experiment  could  decide  which  if  any  of  these  formulae, 
represented  the  structure  of  indigo.  Let  us  try  to  explain  how 
the  organic  chemist  attacks  this  kind  of  a  problem,  for  he  must 
solve  it  beyond  doubt  if  he  hopes  to  synthesize  the  substance. 
By  distillation  of  indigo  with  KOH  one  obtains  aniline  (674),  the 
structure  of  which  is  known  to  be 


The  oxidation  of  indigo  yielded  a  related  substance,  isatine,  for 
which  the  structure  had  been  shown  to  be 


rr 

\/\r 


H 

\ 
C=O 

c/ 


o 

These  facts  indicate  that  the  indigo  molecule  contains  the  atomic 
group 


/-< 


This  group  contains  8  atoms  of  carbon  and  i  of  nitrogen;  since 
indigo  is  Gl6HION2O2,  the  latter  probably  contains  two  isatine 
groups.  Finally,  after  many  years  of  research  the  structure  of 
indigo  was  shown  to  be  the  following: 

H 
/N\/\ 


448  Introduction  to  General  Chemistry 

After  the  structure  of  indigo  had  been  found  it  was  not  long 
until  several  methods  were  devised  by  which  it  could  be  made 
synthetically;  for  once  he  knows  its  structure  the  organic 
chemist  can  attempt  the  building  up  of  a  molecule  with  nearly 
as  great  certainty  of  ultimate  success  as  an  architect  can  con- 
struct a  building  from  prepared  plans.  A  successful  technical 
synthetic  process  must  employ  starting  materials  that  are  suffi- 
ciently abundant  and  cheap.  For  the  artificial  production  of 
indigo  we  start  with  naphthalene  (moth  balls),  CIOH8  (669),  the 
structure  of  which  is 


V 

Oxidation  of  this  substance  yields  phthalic  acid, 

'NCOOH 

OOH 

which  in  turn  yields,  by  the  aid  of  ammonia,  a  substance  having 
the  formula 

r/%H 

\AC/ 

o 

This  by  the  action  of  sodium  hypochlorite  (350)  gives  an  acid 
of  the  following  composition: 


/\/ 


NH> 


The  next  step  is  the  union  of  this  acid  with  chloracetic  acid, 

xv  /NHCH2COOH 

+C1CH2COOH-»|    J^  +HC1 

\/\COOH 


Organic  Compounds  449 

When  the  product  is  heated  with  sodium  hydroxide  it  is  changed 
into  H 

/\/N\ 

(       I  CH2 

\/\c/ 

o 

Upon  oxidation  with  air,  two  molecules  of  this  substance  give 
water  and  a  molecule  of  indigo, 


Artificial  indigo  can  now  be  made  cheaper  than  the  natural  dye 
and  has  nearly  driven  the  latter  from  the  market. 

The  story  of  indigo  is  closely  paralleled  by  that  of  madder, 
an  extract  from  the  root  of  which  dyes  turkey  red.  The  same 
dye  is  now  made  synthetically  from  anthracene,  CI4HIO,  a 
coal-tar  product  resembling  naphthalene. 

The  story  of  indigo  has  been  told  in  some  detail  largely  to 
illustrate  the  method  by  which  a  natural  substance  is  reproduced 
synthetically.  The  steps  in  the  process,  which  in  all  cases  are 
the  same,  are  the  following:  first,  the  isolation  and  purification 
of  the  substance;  second,  the  analysis;  third,  the  molecular- 
weight  determination  to  discover  whether  a  multiple  of  the 
simplest  possible  formula  is  the  true  formula;  fourth,  the  dis- 
covery of  the  structure;  fifth  and  last,  the  synthesis  from  simple 
substances  or  from  the  elements.  Of  these  steps  the  fourth  is 
usually  the  most  difficult,  although  both  this  step  and  the  last 
demand  the  highest  skill  of  the  chemist.  It  would  require  the 
space  of  a  chapter  to  give  even  a  brief  account  of  the  coal-tar 
dyes.  Suffice  it  to  say  here  that  about  nine  hundred  different 
coal-tar  dyes  are  made,  and  of  these  about  three  hundred 
are  in  active  demand.  Each  is  a  definite  chemical  substance, 
the  preparation  of  which  is  on  the  average  as  complicated  a 
process  as  that  of  making  indigo.  At  the  beginning  of  the  war 


450  Introduction  to  General  Chemistry 

90  per  cent  of  all  dyes  used  in  America  were  imported,  largely 
from  Germany.  At  the  close  of  the  war  we  were  making  in 
the  United  States  a  greater  tonnage  of  dyes  than  we  required, 
although  many  less  important  kinds  of  dyes  were  not  yet  made 
in  this  country. 

699.  Synthetic  Medicinals  and  Photographic  Chemicals. — 
At  the  present  time  a  large  number  of  valuable  medicinal  sub- 
stances  are   produced    synthetically   from    coal-tar   products. 
Many  of  these  substances  are  obtained  from  the  same  inter- 
mediate products,  so-called  intermediates,  that  are  used  in 
making  dyes.     Chemicals  used  as  photographic  developers  are 
also  derived  from  similar  intermediates.     These  facts  give  an 
added  significance  to  the  development  of  an  American  dye 
industry,  since  it  means  that  we  shall  also  be  able  to  manufacture 
in  this  country  our  required  medicinals  and  photochemicals.    A 
long  step  in  this  direction  was  taken  during  the  period  of  the  war. 

700.  Rubber. — It  has  long  been  known  that  the  composition 
of  natural  rubber  is  represented  by  the  very  simple  formula 
C5H8,  but  that  the  actual  formula  is  probably  d0Hl6  or  perhaps 
(CIOHI6)».     Thus  rubber  has  the  same  percentage  composition 
as  the  comparatively  cheap  and  abundant  substance  turpentine, 
CIOHl6   (248).     It  is  therefore  not  surprising  that  enormous 
efforts  have  been  made  to  produce  synthetic  rubber.    On  account 
of  the  great  practical  importance  of  this  subject  a  brief  discus- 
sion may  prove  of  interest.     One  of  the  products  of  the  distilla- 
tion of  rubber  is  a  colorless  volatile  liquid  called  isoprene.     This 
has  the  simple  formula  CsHg  and  has  been  shown  to  have  the 

structure 

CH2:C(CH3)CH:CH2 

Isoprene  has  been  made  synthetically  by  several  methods.  It 
is  a  colorless  liquid,  which  by  treatment  with  certain  catalytic 
agents  changes  into  synthetic  rubber: 

2CSH8->CIOHI6. 

As  long  ago  as  1912  a  synthetic-rubber  automobile  tire  that 
had  run  5,000  miles  and  was  still  in  good  condition  was  exhibited 
at  the  International  Congress  of  Applied  Chemistry. 


Organic  Compounds  451 

The  practical  problem  of  the  synthesis  of  rubber  is  then 
resolved  into  the  sufficiently  cheap  production  of  isoprene.  The 
starting  material  must  be  a  compound  of  carbon  and  hydrogen 
with  or  without  other  elements,  which  is  cheap  and  abundant 
and  convertible  with  a  good  yield  into  isoprene.  No  wholly 
satisfactory  technical  method  has  yet  been  found,  for  the  simple 
reason  that  the  cost  of  artificial  rubber  by  any  of  the  processes 
so  far  devised  is  greater  than  that  of  natural  rubber.  Here  then 
is  a  most  interesting  and  important  problem  for  the  future.  To 
the  chemist  who  solves  it  will  come  both  fame  and  wealth. 

701.  Sources  of  Organic  Material. — The  possibilities  of 
synthetic  organic  chemistry  are  very  alluring  to  one  possessed 
of  an  active  imagination;  but  it  is  necessary  to  keep  in  mind 
that  all  synthetic  substances  must  be  made  from  other  carbon 
compounds,  and  that  if  a  substance  is  to  be  made  on  a  practical 
scale  it  can  be  made  only  from  something  sufficiently  abundant 
and  cheap.  These  facts  make  it  worth  while  to  consider  the 
natural  sources  of  organic  materials. 

The  principal  sources  are  (i)  plants,  (2)  animals,  (3)  coal, 
and  (4)  petroleum  and  natural  gas.  It  is  well  known  that  coal 
is  of  vegetable  origin;  and  it  seems  probable  that  petroleum 
comes  from  both  vegetable  and  animal  matter.  Going  back  a 
step  farther  we  recall  that  all  animals  are  dependent  on  plants 
for  nourishment,  so  that  we  find  that  plants  are  the  ultimate 
source  of  all  other  organic  material.  We  have  also  seen  that  plants 
obtain  their  carbon  from  the  carbon  dioxide  of  the  air.  Although 
the  CO2  content  of  the  air  is  but  three  parts  in  10,000  by  volume, 
the  weight  of  this  gas  above  every  square  mile  of  the  earth's 
surface  amounts  to  about  10,000  tons.  Therefore  there  is  in  the 
air  an  immense  reserve  of  the  element  carbon. 

We  have  already  referred  to  the  many  important  constituents 
of  coal  tar  (634,  667).  These  serve  as  the  more  immediate 
starting  materials  for  numerous  organic  syntheses.  The  dis- 
tillation of  wood  also  yields,  in  addition  to  charcoal,  several 
very  useful  products,  of  which  methyl  or  wood  alcohol  (645), 
acetone  (656),  and  acetic  acid  (653)  are  the  most  important. 


CHAPTER  XXVII 
THEORY  OF  DILUTE  SOLUTIONS 

Our  purpose  in  this  chapter  is  to  develop  in  greater  detail 
than  before  the  picture  of  the  condition  of  a  dissolved  sub- 
stance. This  will  be  accomplished  to  best  advantage  by  going 
over  at  the  same  time  parallel  work  done  on  gases. 

702.  The  Reality  of  Molecules. — According  to  the  kinetic 
theory  gases  are  composed  of  tiny  particles  moving  at  random 
with  velocities  of  rifle  bullets  and  spaced  from  each  other  dis- 
tances which  on  the  average  are  large  compared  with  the  diameter 
of  the  particles  (196).     This  picture  has  given  us  an  understand- 
ing of  why  all  gases  diffuse  through  and  occupy  any  space  in 
which  they  are  released.     It  has  also  explained  why  all  gases 
follow  the  same  laws,  namely  those  of   Charles  and  Boyle 
(4,  5).     Further  than  this,  however,  the  kinetic  theory  has 
explained  why  under  great  pressures  and  low  temperature  gases 
do  not  follow  these  laws  closely  (225).     So  great  has  been 
the   assistance   given    by    the    theory    in    understanding    the 
behavior    of    gases    that    the   actual    existence   of   molecules 
has  seemed  to  many  beyond  the  possibility  of  doubt,  even 
though  the  molecules  themselves  have  never  been  seen.     More 
recently,  however,   the    crowning   evidence    in    proof   of    the 
reality  of  molecules  has  been  brought  out  by  work  with  the 
ultra-microscope. 

703.  Perception    of    Molecules. — The    celebrated    French 
physicist  Perrin,  to  some  of  whose  work  on  the  cathode  rays 
we  have  referred  earlier  (478),  says,  "  Direct  perception  of  the 
molecules  in  agitation  is  not  possible  for  the  same  reason  that 
the  motion  of  the  waves  is  not  noticed  by  an  observer  at  too 
great  distance  from  them.    But  if  a  ship  comes  in  sight,  he  is 
able  to  see  that  it  is  rocking,"  and  this  "will  enable  him  to  infer 
the  existence  of  a  possibly  unsuspected  motion  of  the  sea's 
surface."    While  we  cannot  see  molecules  themselves,  we  can 

452 


Theory  of  Dilute  Solutions 


453 


by  devices  now  to  be  described  see  the  phenomenon  which  is 
the  counterpart  of  the  boat  rocking  on  the  invisible  waves. 

In  place  of  a  ship  buffeted  about  by  waves  we  shall  observe 
particles  of  smoke  floating  in  the  air  and  exposed  to  the  bombard- 
ment of  air  molecules.  Although  these  smoke  particles  are 
far  larger  than  the  largest  molecules  (each  consisting  of  many 
molecules),  they  are  still  too  small  to  be  seen  even  under  a 
microscope  unless  they  are  brilliantly  illuminated.  The  smoke 
(from  a  cigarette  or  the  like)  is  contained  in  a  small  glass  box 
beneath  the  objective  of  the  microscope  and  is  illuminated  by  a 
beam  of  light  from  a  projection  lantern  focused  at  a  point  in 


FIG.  98 

the  smoke  just  under  the  objective.  The  arrangement  is  illus- 
trated in  Fig.  98.  The  observer  looks  down  through  the  micro- 
scope upon  tiny  smoke  particles  which  are  illuminated  against 
the  dark  background  by  the  horizontal  rays  from  the  lantern. 
The  field  of  the  microscope  seems  covered  with  a  "milky  way" 
of  bright  stars  which  are  the  flashes  of  light  reflected  to  the 
observer's  eye  by  the  particles.  Thanks  to  the  dark  background 
these  can  be  easily  distinguished. 

The  flat-sided  jar  of  water,  N,  shown  in  the  figure  is  inter- 
posed in  the  beam  from  the  lantern  in  order  to  cool  it  as  much 
as  possible  before  allowing  it  to  strike  the  box;  but  even  so  con- 
vection currents  are  set  up  within  the  box.  As  a  consequence 
of  these  currents  all  the  particles  drift  along  in  one  direction, 
but  if  the  observer  watches  a  single  particle  it  will  be  seen  to  be 
dancing  about  in  a  most  erratic  way.  This  jiggling  motion 
keeps  up  endlessly,  as  long  as  the  particles  themselves  are  in 


454  Introduction  to  General  Chemistry 

suspension.  The  smaller  the  particles  the  more  violent  is  the 
motion.  In  more  refined  apparatus  still  smaller  particles  can 
be  seen;  these  dart  over  distances  many  hundreds  of  times  their 
own  diameter.  The  fact  that  the  motion  is  absolutely  erratic, 
and  that  any  suspension  of  any  substance,  the  individual  particles 
of  which  are  very  small,  exhibits  this  same  motion,  shows  that 
the  latter  is  due  to  the  irregular  bombardment  of  the  particles 
by  the  molecules  of  the  gases  in  which  they  are  suspended.  The 
particles  must  be  small  enough  to  be  moved  appreciably  by  a 
few  more  hits  by  molecules  on  one  side  than  on  the  other.  This 
motion,  called  from  its  discoverer  Brown  the  Brownian  move- 
ment, is  the  phenomenon  which  is  the  counterpart  of  the  rocking 
of  the  boat  upon  the  invisible  waves. 

704.  The  Kinetic  Theory  of  the  Liquid  State. — In  comparing 
liquids  with  gases  we  note  at  once  the  far  greater  density  of  the 
former.  The  kinetic  theory  (198)  describes  molecules  in  the 
liquid  state  as  moving  about,  much  as  they  do  in  the  gaseous 
state,  but  as  greatly  hampered  in  their  motion  on  account  of  the 
crowding  of  the  molecules.  In  liquids  the  average  path  between 
collisions  is  much  shorter  than  in  gases,  so  that  diffusion  through 
the  former  is  very  much  slower  than  through  the  latter.  The 
smaller  distance  between  the  molecules  means  that  the  attrac- 
tion between  them  is  of  greater  moment  than  in  the  gaseous 
state.  This  attraction  falls  off  rapidly  with  increasing  distance, 
so  that  each  molecule  is  affected  by  its  immediate  neighbors 
only.  As  a  result,  within  the  liquid  this  effect  is  balanced,  since 
the  pull  exerted  by  the  neighbors  comes  equally  from  all  sides. 
Only  at  the  surface  of  the  liquid  does  this  force  become  one- 
sided and  so  hampers  the  free  motion  of  the  molecules.  As  a 
result  only  a  few  of  the  faster-moving  molecules  can  escape 
from  the  surface  (199).  According  to  the  kinetic  theory  the 
average  kinetic  energy  of  molecules  in  solution  is  the  same  as 
that  of  molecules  in  the  gaseous  state  at  the  same  temperature, 
so  that  the  light  molecules  in  solution  are  to  be  pictured  as  mov- 
ing very  much  faster  than  the  heavy  ones.  In  many  ways  the 
success  of  the  kinetic  theory  in  helping  us  to  understand  the 
liquid  state  has  been  so  marked  that  its  validity  was  accepted 


Theory  of  Dilute  Solutions  455 

even  before  the  development  of  the  more  direct  evidence  of  the 
ultra-microscope. 

705.  Brownian  Movements  in  Liquids. — The  existence  of 
molecular  motion  in  liquids  can  be  demonstrated  by  the  Brownian 
movements  of  very  tiny  particles  in  suspension  in  liquids,  just 
as  it  can  be  in  the  case  of  smoke  particles.     Suspended  particles 
of  suitable  size,  whether  of  clay,  lampblack,  starch,  or  anything 
else,  all  exhibit  these  movements  irrespective  of  their  composition. 

The  pigment  gamboge,  when  rubbed  in  a  little  water  or 
dissolved  in  alcohol  and  added  to  water,  gives  a  yellowish  turbid 
liquid  which  consists  of  suspended  droplets  of  suitable  size  to 
show  the  B  rownian  movements.  A  microscope  with  oil-immersion 
objective  is  best  used.  If  a  drop  of  the  gamboge  suspension  is 
mounted  for  observation  under  the  objective  and  viewed  with 
the  ordinary  illumination  reflected  from  the  mirror  below  the 
stage  of  the  instrument  many  tiny  particles  will  be  seen  in  the 
field.  Careful  observation  will  show  that  the  smaller  particles 
are  in  a  constant  tremor,  moving  erratically  over  paths  approxi- 
mately one  to  several  times  their  own  diameters  in  length.  It 
seems  difficult  for  the  observer  to  believe  that  these  moving 
particles  are  not  alive,  so  lifelike  is  their  motion! 

706.  The  Revelations  of  the  Ultra-Microscope. — Suspensions 
of  far  smaller  particles  may  be  observed  if  instead  of  the  ordinary 
microscope  we  use  the  ultra-microscope.    To  build  a  crude  form 
of  the  latter  we  may  employ  the  same  set-up  as  was  used  for 
the  examination  of  the  smoke  particles,  except  that  we  must 
use  an  oil-immersion  lens  and  a  so-called  condenser.     In  this 
case  the  light  from  the  lantern  is  again  made  to  pass  through 
water,  but  this  time,  instead  of  being  directed  next  against  the 
object,  it  is  made  to  strike  the  mirror  below  the  stage  of  the 
microscope  and  is  reflected  into  the  condenser,  which  directs  it 
obliquely  against  the  droplet  on  the  slide,  as  shown  by  the  path 
of  the  arrows  in  Fig.  99.     A  drop  of  pure  water  placed  on  the 
slide  appears  only  as  a  dark  field  to  the  observer;  but  if  there 
are  tiny  particles  present  in  it  flashes  of  light  are  reflected  from 
them  to  the  eye.    Zsigmondy,  who  with  Siedentopf  perfected 
the  ultra-microscope,  describes  his  first  view  of  a  fine  suspension 


456  Introduction  to  General  Chemistry 

of  gold  particles  in  water  thus:  "The  small  gold  particles  move 
and  that  with  astonishing  rapidity.     A  swarm  of  gnats  in  a 
sunbeam  will  give  one  an  idea  of  the  motion  of  the  gold  particles. 
....  They  hop,  dance,  jump,  dash  together  and  fly  away  from 
each  other  so  that  it  is  difficult  to  get  one's  bearings.''    The  very 
fine  particles  seen  under  the  ultra-microscope  are  observed  to 
travel  much  more  rapidly  and  to  cover 
»\  /,  greater  distances  than  those  previously  de- 

scribed. It  should  be  realized  that  any 
suspension  in  any  liquid  shows  this  same 
erratic  motion.  Specimens  have  been  under 
observation  for  years,  but  no  abatement 
of  the  Brownian  movement  has  been  ob- 
served. Suspensions  in  liquids  taken  from 
FIG.  99  the  interior  of  quartz  crystals  show  these 

same  motions,  and  yet  they  are  thousands 
of  years  old !  We  have,  therefore,  the  same  type  of  evidence  of 
the  reality  of  the  existence  of  swift-moving  molecules  in  liquids 
as  we  have  in  gases. 

707.  Visible  Atmospheres. — Perrin  was  able  to  show  that 
fine  suspensions  may  be  considered  as  visible  atmospheres.  The 
pressure  due  to  the  impacts  of  the  ultra-microscopically  visible 
particles  is  proportional  to  the  concentration  and  to  the  absolute 
temperature,  just  as  the  pressure  due  to  gas  molecules  is  pro- 
portional to  their  concentration  and  to  the  absolute  temperature. 
The  actual  number  (194)  of  molecules  in  a  gram  molecular 
weight  of  a  substance  in  the  gaseous  state  has  been  estimated 
as  6  X  io23.  Perrin  estimated  the  number  of  suspension  particles 
in  what  would  be  a  gram  molecular  weight  of  a  suspended  sub- 
stance, namely  the  weight  of  substance  in  22  .4  liters  of  a  suspen- 
sion the  particles  of  which  at  o°  exert  a  pressure  of  760  mm. 
The  numbers  found  for  a  great  variety  of  suspensions  were 
close  to  6  X  io23.  The  viscosity  of  the  suspension  medium  seemed 
to  make  no  difference  in  the  result.  It  did  not  matter  whether 
the  suspension  medium  was  pure  water  or  the  very  viscous 
glycerine.  The  average  weight  of  the  particles  was  varied 
70,000  fold,  and  yet  the  results  were  the  same.  Thus  the  gas 


Theory  of  Dilute  Solutions 


457 


laws  and  Avogadro's  hypothesis  apply  to  these  suspensions. 
We  are  then  to  think  of  fine  suspensions  as  miniature  atmospheres 
of  colossal  molecules  which  are  actually  visible. 

708.  Particles  in  Solution  vs.  Particles  in   Suspension.— 
Particles  visible  in  the  ultra-molecular  microscope  are  called 
suspensions,  but  particles  just  too  small  to  be  seen  in  this  way 
must  be  counted  as  in  solution.     The  change  from  solution  to 
suspension  should  be  thought  of  as  gradual  rather  than  abrupt, 
for  the  distinction  seems  to  be  fundamentally  one  of  the  size  of  the 
particle. 

709.  Substances  in  Solution  vs.  Gases. — The  knowledge  that 
particles  in  suspension  in  a  liquid  act  like  huge  gas  molecules 
inclosed  in  the  same  volume  as  that  of  the  solution  leads  at  once 
to  the  conception  that  substances  in  solution  may  also  be  con- 
sidered in  the  same  light.     To  test  this  theory  the  pressure  of 
the  dissolved  molecules  must  be  measured  separately  from  that 
of  the  molecules  of  the  solvent.     This  is  obviously  not  a  simple 
undertaking,  since  the  full  impacts  of  molecules  in  solution  are 
not  felt  by  a  surface  in  contact  with  the  latter 

because  of  the  attraction  inward  of  the  molecules 
which  come  to  the  surface  layer.  The  same  force 
which  prevents  the  free  migration  of  molecules 
from  the  surface  of  the  liquid  prevents  the 
measurement  of  the  force  of  their  impacts  by  such 
devices  as  are  used  to  measure  gas  pressure. 
However,  devices  to  suit  the  purpose  have  been 
found,  and  these  will  now  be  studied.  Since 
they  depend  on  the  operation  of  a  semipermeable 
membrane,  a  simple  example  of  the  latter  will  first 
be  described. 

710.  Semipermeable  Membranes. — A  so- 
called  semipermeable  membrane  is  a  membrane 
which  is  permeable  to  one  of  two  substances  in 
a  mixture  but  not  to  the  other.     A  layer  of 

thin  cloth  (muslin)  stretched  over  the  tubulated  end  of  the  oil 
manometer,  as  shown  in  Fig.  100,  and  wet  with  water  will  serve 
as  a  membrane  permeable  to  ammonia  but  not  to  air.  The 


NH3 


FIG.  100 


458 


Introduction  to  General  Chemistry 


stopper  at  A  permits  the  insertion  of  red  litmus  paper  into  the 
inner  chamber,  as  shown  in  the  figure.  When  ammonia  is  col- 
lected in  a  beaker  inverted  over  the  membrane,  the  oil  in  the 
manometer  moves,  showing  an  increase  of  pressure  in  the  inner 
chamber.  Meantime  the  litmus  paper  is  seen  to  turn  blue,  thus 
showing  that  ammonia  has  entered.  This  it  has  done  by  dis- 
solving in  the  upper  surface  of  the  water  film  and  diffusing 
through  it.  As  a  result  some  ammonia  molecules  soon  appear 
at  the  inner  surface  of  the  film  and  then 
escape  into  the  air  within  the  apparatus,  just 
as  in  the  case  of  the  diffusion  through  a  gas 
(191)  the  migration  of  ammonia  molecules 
through  the  water  and  air  layers  proceeds 
spontaneously  from  regions  of  higher  concen- 
tration to  those  of  lower  concentration  until 
there  is  no  longer  any  difference  between  the 
two.  Meantime  the  air  inclosed  in  the  ap- 
paratus is  not  able  to  escape,  since  it  is  not 
appreciably  soluble  in  water,  and  therefore 
the  total  pressure  increases. 

711.  The  Solution  Cell. — A  similar  appa- 
ratus for  use  with  solutions  may  be  easily 
prepared  for  demonstration  purposes  by  at- 
taching a  parchment-paper  thimble  to  a  glass 
tube,  as  shown  in  Fig.  101.  The  thimble  is 
filled  with  concentrated  sugar  solution  colored 
with  cochineal  and  immersed  in  a  beaker  of 
pure  water.  Water  can  pass  through  the 
parchment  wall  in  either  direction,  but  the 
sugar  and  the  cochineal  cannot  do  so  appre- 
3  ciably.  Since  the  color  can  only  accompany 
FIG.  1 01  the  sugar,  it  serves  as  a  convenient  marker 

for  the  presence  of  the  latter.  Soon  after 
the  cell  is  filled  and  the  stopcock  of  the  filling  tube,  A,  is 
closed  the  level  of  the  sugar  solution  begins  to  rise  in  the  verti- 
cal tube.  In  a  few  hours  the  surface  of  the  sugar  solution  will 
be  three  or  four  feet  above  the  level  of  the  water  in  the  outside 


Theory  of  Dilute  Solutions 


459 


beaker.  There  is  obviously  a  close  analogy  between  this  experi- 
ment and  the  preceding  one.  The  water  in  the  former  cor- 
responds to  the  ammonia  in  the  latter,  since  each  can  pass 
through  the  membrane.  The  sugar,  like  the  air,  cannot  pass 
through  the  membrane.  Whatever  the  mechanism  by  which 
the  cell  operates,  it  is  plain  that  the  weight  of  the  column  of 
solution  is  borne  by  the  membrane,  except  for  the  slight  assist- 
ance given  by  the  pressure  exerted  against  it  by  the  water  in  the 
outside  beaker.  If  the  column  is  allowed  to  get  too  high  the 
thimble  will  be  broken.  For  accurate  measurements  stronger 
membranes  must  be  used,  and  they  must  also  be*  of  substances 
more  truly  impermeable  to  sugar.  The  best  membranes  have  been 
made  by  precipitating  copper  ferrocyanide 
in  the  pores  of  unglazed  porcelain  cups. 

A  cell  of  this  type  is  shown  in  Fig.  102. 
The  porcelain  cup  which  holds  the  sugar 
solution  is  fitted  with  a  mercury  manom- 
eter, the  closed  end  of  which  contains  nitro- 
gen gas.  As  water  enters  the  cell  through 
the  copper  ferrocyanide  membrane  deposited 
in  the  pores  of  the  cup,  the  nitrogen  in  the 
manometer  is  compressed.  Since  the  manom- 
eter is  very  narrow,  only  a  small  increase  of 
volume  on  the  part  of  the  solution  in  the  cell 
results  in  a  very  great  compression  of  the 
nitrogen.  The  change  in  volume  of  the  latter 
(together  with  the  very  small  difference  in 
levels  of  the  mercury  in  the  manometer  arms) 
gives  the  necessary  data  for  calculating  the 
pressure  within  the  cup.  When  no  further 
change  in  volume  of  the  nitrogen  is  seen,  the  pressure  of  the 
latter  is  just  sufficient  to  check  the  inflow  of  water  through  the 
membrane.  The  pouring  of  the  water  into  the  sugar  solution 
through  the  membrane  is  called  osmosis,  and  the  pressure  which 
checks  this  flow  is  called  the  osmotic  pressure  of  the  solution. 

712.  Results   of    Osmotic-Pressure    Measurement. — Table 
XXXIII  gives  the  osmotic  pressures  of  cane-sugar  solutions 


FIG.  102 


460 


Introduction  to  General  Chemistry 


at  60°  at  the  concentrations  indicated.  From  indirect  evidence 
chemists  long  ago  made  up  their  minds  that  the  molecular 
weight  of  sugar  is  342,  corresponding  to  the  formula  CI2H22OIX 
(682).  Hence  we  may  give  the  concentration  in  percentage  by 
weight  of  sugar  in  solution  and  also  in  terms  of  gram  molecular 
weights  of  sugar  to  1,000  grams  of  water,  i.e.,  weight-molar  con- 
centration. The  osmotic  pressure  is  given  first  in  Ibs.  per  square 
inch,  in  order  that  the  magnitude  of  these  pressures  may  be 
visualized,  and  then  in  terms  of  atmospheres.  From  the  table 
it  is  evident  that  even  for  a  3  per  cent  solution  the  membrane 

TABLE  XXXIII 

OSMOTIC  PRESSURES  OF  SUGAR  SOLUTIONS 


C  ONCENTRATION 

OSMOTIC  PRESSURE 

GAS  PRESSURE 

IN 

ATMOSPHERES 

RATIO 
O.P.  TO  G.P. 

Percentage 

Weight-Molar 

Lbs.PerSq.  In. 

Atmospheres 

3-3 

O.  I 

40 

2.72 

2.72 

.00 

6-4 

.  2 

80 

5-44 

5-43 

.00 

9-3 

-3 

1  2O 

8.14 

8.15 

.00 

12.0 

•  4 

1  60 

10.87 

10.87 

.00 

I4.6 

•5 

2O2 

13.67 

I3-58 

.01 

I7.I 

.6 

244 

16.54 

16-30 

.01 

J9-3 

•7 

285 

19.40 

19.02 

.02 

21-5 

0.8 

329 

22.33 

21-73 

•03 

must  stand  a  pressure  of  40  Ib.  to  the  square  inch.  For  a 
20  per  cent  solution  the  membrane  must  stand  a  pressure  of 
over  300  Ib.  per  square  inch.  Students  will  readily  appreciate 
the  experimental  skill  necessary  to  make  these  measurements, 
when  they  think  of  their  own  difficulties  in  securing  an  apparatus 
tight  enough  to  stand  less  than  a  pound  of  pressure  without 
leaking !  In  the  fifth  column  is  given  the  pressure  which  a  like 
molecular  concentration  of  hydrogen  would  exert  at  the  same 
temperature.  The  close  agreement  between  osmotic  pressure 
and  gas  pressure  is  at  once  apparent  from  a  comparison  of  the 
fourth  and  fifth  columns.  This  is  brought  out  in  the  sixth 
column,  which  represents  the  ratio  of  osmotic  pressure  to  gas 
pressure.  Where  the  ratio  is  equal  to  i .  oo,  the  two  may  be 
regarded  as  equal.  As  has  been  said,  this  data  is  for  solutions 


Theory  of  Dilute  Solutions 


461 


at  60°.  How  osmotic  pressure  compares  with  gas  pressure  at 
other  temperatures  is  shown  in  Table  XXXIV,  which  represents 
the  ratio  of  osmotic  pressure  to  gas  pressure  for  solutions  at 
other  temperatures  as  well.  The  data  for  both  tables  have  been 
adapted  from  the  very  fine  work  of  Morse  and  Frazer,  of  Johns 
Hopkins  University. 

In  both  Table  XXXIII  and  Table  XXXIV  it  is  apparent 
that  the  osmotic  pressures  are  indeed  close  to  the  corresponding 
gas  pressures.  If  we  look  at  the  deviations  from  the  gas  laws 
by  noting  how  much  greater  than  i .  oo  the  ratios  are,  we  find 


TABLE  XXXIV 
RATIO  OF  OSMOTIC  PRESSURE  TO  GAS  PRESSURE  AT  VARIOUS  TEMPERATURES 


Concentration 

0° 

20° 

40° 

60° 

O    I 

( 

II) 

08 

OO 

OO 

O   ^ 

06 

06 

O2 

oo 

o  t; 

07 

O7 

05 

OI 

o  8 

.OQ 

.OQ 

O7 

.03 

I  .0  

.  12 

.  12 

.09 

.04 

that  they  are  larger  the  lower  the  temperature  and  the  more 
concentrated  the  solution.  This  reminds  one  of  the  deviations 
of  the  behavior  of  the  gases  themselves  from  strict  agreement 
with  the  gas  laws  (225). 

713.  Interpretation  of  Osmotic-Pressure  Measurements. — 
How  are  we  to  explain  the  striking  results  of  osmotic-pressure 
measurements?  The  fact  that  the  measured  osmotic  pressure 
of  a  solution  equals  the  pressure  due  to  a  gas  at  like  temperature 
and  of  like  molecular  concentration  as  that  of  the  dissolved 
substance  in  question  is  a  coincidence  so  remarkable  that  we 
cannot  call  the  agreement  one  of  chance.  Apparently  the  cell 
operates  because  water  is  at  a  different  concentration  on  the 
two  sides  of  the  membrane  just  as  the  air-ammonia  cell  operated 
because  ammonia  which  could  pass  through  the  water  mem- 
brane was  at  first  more  concentrated  on  one  side  than  on  the 
other.  In  the  sugar-water  cell  presumably  the  pressure  due 
to  the  impacts  of  the  molecules  against  the  membrane  is  at  first 


462  Introduction  to  General  Chemistry 

the  same  from  both  sides;  hence  the  pressure  of  water  molecules 
inside  the  cell  is  short  of  that  outside  by  just  the  pressure  due 
to  the  impacts  of  the  sugar  molecules.  This  pressure  must  be 
put  upon  the  solution  in  order  that  the  water  may  be  made 
to  pass  back  from  solution  to  pure  water  at  the  same  rate  that 
it  is  coming  into  the  cell  and  equilibrium  thus  be  established. 
In  this  way  we  may  explain  the  operation  of  the  cell  and  the 
peculiar  values  found. 

In  our  previous  work  we  have  always  given  concentrations 
of  solutions  in  terms  of  weight  of  dissolved  substance  per  liter  of 
solution.  In  the  above-mentioned  work  concentration  is  given 
in  terms  of  weight  of  dissolved  substance  per  liter  of  solvent, 
since  the  values  so  obtained  are  a  little  nearer  to  the  correspond- 
ing gas  pressures,  presumably  because  the  amount  of  solvent 
is  a  measure  of  the  actual  free  space  through  which  the  dissolved 
particles  move. 

In  measuring  the  pressure  which  must  be  put  upon  a  solu- 
tion to  check  the  inflow  of  solvent  into  it  when  separated  from 
the  solvent  by  a  semipermeable  membrane,  we  are  measuring 
a  force  which  is  equal  to  the  pressure  exerted  by  the  impacts 
of  the  molecules  of  the  dissolved  substances.  From  such 
measurements  we  learn  that  substances  in  dilute  solution  act 
much  as  though  they  were  in  the  gaseous  state,  at  the  same 
temperature  and  inclosed  in  volumes  equal  to  the  volume  of  the 
pure  solvent  in  the  solution. 

714.  Other  Theories  of  Osmosis.— The  theory  just  presented 
of  the  mechanism  by  which  the  semipermeable  membrane  cell 
operates  is  not  the  only  possible  explanation  of  the  facts.  For 
instance,  some  prefer  to  regard  the  motion  of  the  solvent  through 
the  membrane  as  due  to  an  attraction  by  the  solution.  Osmotic 
pressure  would  be  the  result  of  a  sort  of  suction.  As  yet  the 
theory  presented  above  has  been  by  far  the  more  useful. 

715-  Van't  Hoff's  Theory  of  Solutions.— The  great  Dutch 
chemist  van't  Hoff  was  the  first  to  point  out  that  the  gas  laws 
should  hold  for  substances  in  very  dilute  solution.  He  drew 
his  conclusions  from  considerations  of  a  purely  theoretical 
nature  at  a  time  (1887)  when  the  data  at  hand  for  the  experi- 


Theory  of  Dilute  Solutions  463 

mental  confirmation  of  his  theory  were  far  less  accurate  than 
those  given  in  Tables  XXXIII  and  XXXIV  and  showed  far  less 
clearly  the  truth  of  his  statements.  Van't  Hoff  pointed  out 
that  the  laws  of  gases  should  be  followed  by  substances  in  very 
dilute  solution  only,  since  in  the  more  concentrated  solutions 
the  formation  of  complexes  between  the  solvent  and  the  dis- 
solved substance  and  between  the  molecules  of  the  latter  make 
conditions  which  are  not  parallel  to  those  found  in  gases  which 
follow  the  simple  gas  laws. 

716.  The  Avogadro-van't  Hoff  Hypothesis. — Van't  Hoff  also 
asserted  the  so-called  Avogadro-van't  Hoff  hypothesis,  a  generali- 
zation of  tremendous  importance  in  the  growth  of  modern  chem- 
istry.    This  hypothesis  is  simply  the  application  of  Avogadro's 
hypothesis  to  substances  in  solution.     Equal  volumes  of  solu- 
tions of  all  substances  which  have  equal  osmotic  pressures  at  the 
same  temperature  contain  the  same  number  of  dissolved  molecules. 

717.  Molecular   Weights  from   Osmotic-Pressure    Data.— 
As  a  first  result  of  the  establishment  of  this  principle  we  have 
learned  how  to  find  the  relative  weights  of  molecules  in  solution 
(molecular  weights)  by  finding  what  weight  of  the  substance  in 
question  will  exert  76  cm.  of  osmotic  pressure  at  o°  when  dis- 
solved in  22 .4  liters  of  solvent. 

718.  Raoult's    Work.     Shorter    Methods    of    Determining 
Molecular  Weights. — Before  van't  Hoff's  work  was  published, 
Raoult,  a  French  chemist,  working  with  a  great  variety  of  sub- 
stances the  molecular  weights  of  which  were  known,  had  shown 
that  a  gram  molecular  weight  of  each  dissolved  in  a  given  weight 
of  solvent  produced  about  the  same  lowering  of  the  vapor 
pressure  of  the  latter,  provided  the  dissolved  substance  was  not 
volatile.     Van't  Hoff  pointed  out  (on  the  basis  of  reasoning 
which  we  need  not  now  consider)  that  this  is  a  necessary  con- 
sequence of  the  fact  that  Avogadro's  hypothesis  applied  to  dilute 
solutions.     Therefore   the   molecular   weight   of   an   unknown 
substance  in  solution  may  be  determined  by  finding  what  weight 
of  this  substance  is  required  to  give  a  solution  of  vapor  pressure 
equal  to  that  of  a  half-molar  solution  of  a  substance  the  molec- 
cular  weight  of  which  is  known.     Twice  the  value  found  would 


464       .  Introduction  to  General  Chemistry 

be  the  molecular  weight  of  the  new  substance.  Raoult's  observa- 
tions are  for  dilute  solutions,  so  that  the  work  should  not  be 
done  with  concentrations  greater  than  half -molar. 

We  have  already  seen  that  one  result  of  the  lowering  of  the 
vapor  pressure  of  a  liquid  by  dissolving  a  non- volatile  foreign 
substance  in  it  is  that  the  solution  boils  at  a  higher  temperature 
than  the  pure  solvent  (128).  Raoult  showed  that  the  elevation 
of  the  boiling-point  is  proportional  to  the  number  of  gram  molec- 
ular weights  of  substance  dissolved  in  unit  weight  of  solvent.  If 
one  gram  molecular  weight  of  a  non- volatile  substance  is  dis- 
solved in  i,ooog.  of  water  the  boiling-point  is  raised  0.50°. 
This  is  the  so-called  molar  elevation  for  water.  A  gram  molec- 
ular weight  of  a  substance  dissolved  in  1,000  g.  of  ether  elevates 
the  boiling-point  of  the  latter  2 . 20°.  Here  then  is  another  way 
of  determining  molecular  weights  of  substances  in  solution.  We 
have  only  to  find  how  much  of  the  new  substance  must  be  dis- 
solved in  i,ooog.  of  solvent  to  produce  the  molar  elevation  of 
the  boiling-point,  in  order  to  find  the  molecular  weight  of  the 
substance  in  question. 

The  freezing-point  of  a  solution  affords  still  another  method. 
The  depression  of  the  freezing-point  of  the  solvent,  like  the  elevation 
of  the  boiling-point,  is  proportional  to  the  number  of  gram  molec- 
ular weights  present  per  unit  weight  of  solvent,  provided  the  latter 
is  the  only  substance  which  separates  from  solution  during  the 
freezing.  Water  solutions  which  contain  one  gram  molecular 
weight  of  dissolved  substance  per  i,ooog.  of  water  freeze  at 
-1.85°.  Hence  the  molar  depression  of  the  freezing-point 
of  water  may  be  taken  as  i  .85°.  That  of  the  solvent  benzene 
(667)  is  5.3°.  Accordingly  we  need  only  find  what  weight  of 
substance  is  necessary  to  produce  the  molar  depression  of  the 
freezing-point  of  a  suitable  solvent  to  find  the  molecular  weight 
of  the  substance  in  question  in  the  solvent  used. 

719.  The  Importance  of  the  Work  of  van't  Hoff  and  of 
Raoult. — The  general  acceptance  of  Avogadro's  hypothesis 
(about  1860)  initiated  a  new  era  in  chemistry.  A  second  great 
development  of  chemistry  began  in  1887,  with  the  announcement 
of  the  Avogadro  van't  Hoff  hypothesis.  The  latter  furnished 


Theory  of  Dilute  Solutions  465 

the  basis  for  the  understanding  of  reactions  in  dilute  solution, 
just  as  the  original  hypothesis  had  done  for  gas  reactions.  Much 
of  the  rapidity  of  the  development  of  the  work  on  dilute  solutions 
is  due  to  the  relatively  simple  experimental  methods  of  Raoult 
for  determining  molecular  weights  of  dissolved  substances.  The 
immediate  result  of  the  work  of  these  two  men  was  the  presenta- 
tion of  new  and  independent  evidence  concerning  the  ionic  hy- 
pothesis which  was  just  being  developed  by  Arrhenius  (in  1887), 
mainly  from  work  on  conductivity  measurements. 

720.  Molecular  Weights  of  Acids,  Bases,  and  Salts. — When 
the  molecular  weights  of  acids,  bases,  and  salts  were  determined 
in  water  solution  abnormal  results  were  found.    Thus  in  the 
case  of  hydrogen  chloride,  which  as  a  gas  has  the  molecular 
weight  36.5,  the  molecular  weight  in  water  solution  was  found 
to  be  much  less  than  this,  and  furthermore,  when  measurements 
were  made  on  solutions  of  greater  and  greater  dilution,  the 
results  found  grew  smaller,  reaching  about  18 .3  as  a  limit. 

The  significance  of  abnormally  low  molecular  weight  in  the 
case  of  ammonium  chloride  vapor  (529)  was  shown  to  be  due 
to  dissociation  into  ammonia  and  hydrogen  chloride.  Arrhenius 
was  quick  to  point  out  that  the  abnormally  low  molecular 
weights  of  electrolytes  must  also  be  due  to  the  fact  that  the 
original  molecules  must  have  dissociated  into  ions  in  water 
solution.  The  fact  that  these  molecular  weights  get  smaller 
with  increasing  dilution  is  evidence  of  increased  dissociation 
with  dilution;  but  finally,  when  the  degrees  of  dissociation 
were  calculated  from  freezing-point  data  and  compared  with 
the  values  found  for  the  same  concentrations  by  the  conductivity 
method  (408),  the  two  were  found  to  agree  so  closely  that  little 
doubt  remained  of  the  correctness  of  the  ionic  hypothesis. 

721.  The  Degree  of  Dissociation  from  Molecular- Weight 
Determinations. — The  method  of  calculation  of  the  degree  of 
dissociation  from  molecular-weight  determinations  is  best  illus- 
trated by  an  example.    In  the  case  of  potassium  chloride  the 
simplest  formula  is  KC1,  indicating  a  molecular  weight  of  74.6. 
If  KC1  ionizes  thus, 


466  Introduction  to  General  Chemistry 

it  is  plain  that  if  100  molecules  were  originally  present  and 
80  per  cent  of  them  broke  down  into  ions,  the  total  number  of 
free  particles  present  would  be  180  instead  of  100.  If  we  start 
with  74. 6  g.  of  KC1  in  i,ooog.  of  water  and  80  per  cent  is 
dissociated,  we  shall  have  present  0.2  gram  molecular  weight 
of  KC1,  0.8  gram  molecular  weight  of  potassium  ion,  and 
0.8  gram  molecular  weight  of  chlorine  ion,  making  a  total  of 
i .  8  gram  molecular  weights  from  the  74 . 6  g.  of  salt  dissolved 
instead  of  one  gram  molecular  weight.  The  observed  depres- 
sion of  the  freezing-point  of  water  will  be  i .  8  times  as  great  as 
is  expected  for  one  gram  molecular  weight  of  total  substance 
dissolved  in  the  same  volume.  Hence  the  apparent  molecular 
weight  of  KC1  calculated  from  the  freezing-point  of  a  solution 
80  per  cent  ionized  will  be  74.6-1-1.8  =  41.4.  Conversely,  if 
the  molecular  weight  of  a  KC1  solution  found  in  the  usual  way 
by  the  freezing-point  method  is  41 .4,  the  degree  of  dissociation 
of  the  solution  is  found  thus:  74.6-1-41.4  =  1.80.  From  this 
we  conclude  that  the  dissociation  is  80  per  cent. 

722.  Degree  of  Dissociation  of  Electrolytes  by  Two 
Methods. — The  degrees  of  dissociation  of  a  number  of  electro- 
lytes at  different  concentrations  as  calculated  from  molecular- 
weight  determinations  by  the  method  just  illustrated  are  shown 
in  Table  XXXV  (line  I  for  each  substance)  in  comparison  with 
like  data  for  the  same  substance  as  found  from  electrical  con- 
ductivities (line  II).  (In  the  more  concentrated  solutions  the 
latter  data  have  been  corrected  for  changes  in  the  viscosity  of 
the  solution  with  concentration.  Obviously,  if  the  viscosity  of 
the  solution  increases,  the  rapidity  with  which  the  same  ions  will 
move  under  the  same  attractive  force  will  be  diminished  so  that 
a  correction  is  necessary.)  Concentrations  are  given  in  terms  of 
molecular  weights  per  liter  except  in  the  cases  of  salts  having 
bivalent  ions.  In  these  cases  the  fraction  J  indicates  that  the 
concentration  is  one-half  the  concentration  indicated  at  the  head 
of  the  column.  The  solutions  represented  in  the  same  column 
contain  equivalent  weights  of  materials  (403). 

Apparently  the  agreement  between  the  results  of  the  two 
methods  is  within  a  few  per  cent  for  concentrations  up  to 


Theory  of  Dilute  Solutions 


467 


0.5  molar  for  salts  consisting  of  two  univalent  ions  in  com- 
bination, and  for  concentrations  up  to  0.25  molar  for  salts 
consisting  of  two  univalent  ions  united  to  one  bivalent  ion. 
With  increasing  concentrations  for  both  types  the  divergences 
become  larger.  Data  for  salts  made  up  of  higher  valent  ions 
show  even  greater  disagreements  in  the  two  methods.  However, 

TABLE  XXXV 

DEGREES  OF  DISSOCIATION* 


Concentrations 

O.OI 

0.05 

0.  10 

0.  20 

0.50 

KC1I.. 
KC1II 

94% 

04 

89% 
80 

86% 
86 

83% 

8* 

80% 
78 

NaClI 

04 

80 

88 

81 

82 

NaCl  II  .    . 

04 

88 

8<c 

82 

77 

KBrl.. 

OO 

86 

84 

81 

KBrll  

04 

80 

86 

8? 

77 

HC1I.  . 

97 

93 

HC1II  

97 

94 

HNO3I 

06 

QI 

HNO3  II  . 

07 

04 

BaCl2(£)  I.  . 

88 

82 

70 

76 

BaCl2(|)  II  

88 

80 

76 

72 

CaCl2(£)  I.  . 

84 

82 

80 

CaCl2Q)  II  

80 

76 

77 

Pb(NO3)2(|)  I 

8? 

72 

6? 

cry 

43 

Pb(NO3)2G)  II. 

8s 

71 

64 

S6 

AC 

K2SO4(£)  I.. 

QO 

78 

72 

67 

^7 

K2SO4(|)  II  

87 

77 

72 

67 

62 

*  Adapted  from  the  work  of  Noyes  and  Falk. 

the  complications  which  arise  under  both  these  circumstances 
make  these  discrepancies  appear  to  be  in  keeping  with  the  funda- 
mental theory. 

723.  Summary. — The  material  developed  in  this  chapter 
may  be  summarized  as  follows:  The  existence  of  free-moving 
molecules  in  liquids  as  described  by  the  kinetic  theory  has  been 
put  beyond  the  possibility  of  doubt  by  the  discovery  that  all 
very  fine  suspensions  are  in  constant  motion.  Perrin  has  shown 


468  Introduction  to  General  Chemistry 

that  these  fine  suspensions  behave  like  atmospheres  of  huge 
molecules.  The  kinetic  energy  of  a  suspension  particle  is  equal 
to  that  of  a  gas  molecule  at  the  same  temperature.  In  fact, 
the  gas  laws  may  be  applied  to  these  suspensions.  Even  when 
the  average  mass  of  the  molecules  is  varied  70,000  fold,  the 
same  results  are  found.  Since  the  distinction  between  an 
extremely  fine  suspension  and  a  solution  is  fundamentally  one 
of  the  size  of  the  particles,  it  seems  plausible  that  a  substance  in 
dilute  solution  might  also  be  found  to  behave  as  though  in  the 
gaseous  state. 

The  pressure  exerted  by  the  molecules  of  a  dissolved  substance 
is  determined  by  measuring  an  equal  and  opposite  pressure, 
namely  that  of  the  inflow  of  solvent  into  the  solution  through 
a  membrane  permeable  to  the  solvent  but  not  to  the  dissolved 
substance.  The  pressures  so  measured  are  found  to  be  very 
nearly  equal  to  the  pressure  which  a  gas  would  exert  if  of  the 
same  molecular  concentration  and  at  the  same  temperature  as 
the  dissolved  substance.  Thus  the  analogy  between  substances 
.in  dilute  solution  and  in  the  gaseous  state  is  completely  borne 
out.  Van't  Hoff,  who  was  the  first  to  develop  the  theory  of 
dilute  solutions,  pointed  out  that  Avogadro's  Law  should  hold 
for  substances  in  dilute  solution.  On  the  basis  of  this  assump- 
tion the  molecular  weight  of  a  substance  in  solution  can  be 
obtained  by  finding  what  weight  of  the  latter  dissolved  in  22 .4 
liters  will  exert  760  mm.  osmotic  pressure  at  o°.  Raoult,  work- 
ing with  a  great  variety  of  substances,  the  molecular  weights  of 
which  were  known,  showed  that  a  gram  molecular  weight  of 
each  produced  about  the  same  change  in  the  vapor  pressure  of 
the  solvent.  The  corresponding  changes  in  the  boiling-point 
and  freezing-point  can  be  easily  measured.  The  theoretical 
deductions  of  Van't  Hoff  showed  that  the  experimental  methods 
of  Raoult  were  simply  other  methods  for  determining  the  rela- 
tive weights  of  the  molecules  as  they  existed  in  the  solutions. 

On  the  basis  of  the  work  of  these  two  men  the  existence  of 
particles  smaller  in  weight  than  the  original  molecules  of  elec- 
trolyte dissolved  in  water  solutions  was  established.  Further- 
more the  amounts  of  these  smaller  particles  were  found  to  be 


Theory  of  Dilute  Solutions  469 

in  agreement  with  the  requirements  of  the  ionic  theory.  As  a 
consequence  the  discovery  of  these  facts  put  the  ionic  theory 
on  a  firm  foundation. 

724.  Osmosis  in  Nature. — Osmosis  is  one  of  the  important 
processes  operative  in  nature.  This  is  illustrated  by  the  effect 
of  water  on  wilted  plants.  The  latter  have  lost  water  in  the 
air,  but  upon  being  put  into  water  they  " freshen"  because  the 
water  pours  through  the  cell  walls  and  thus  puts  a  pressure 
against  the  latter  which  stiffens  them.  This  pressure  could  not 
accumulate  if  the  cell  walls  were  not  only  membranes,  permeable 
to  water,  but  also  impermeable  to  much  of  the  other  contents 
of  the  cell.  If  a  fresh  flower  is  put  into  concentrated  sugar 
solution,  in  a  very  few  minutes  it  will  be  found  to  be  badly 
wilted.  This  occurs  because  the  concentration  of  water  is  less 
in  the  sugar  solution  than  it  is  in  the  plant  cells  (713).  The 
effect  of  sugar  in  withdrawing  the  juice  from  fruit  is  another  well- 
known  phenomenon  of  the  same  type.  The  foregoing  are  easily 
observed  examples.  However,  many  less  easily  observed  but 
important  physiological  processes  both  of  plants  and  animals 
depend  on  osmosis 


CHAPTER  XXVIII 
DISPERSE  SYSTEMS 

725.  Introduction. — The  surface  layer  of  a  massive  solid  or 
of  a  liquid,  as  it  is  contained  in  an  ordinary  vessel,  is  of  negligible 
importance  chemically  compared  to  its  total  mass.     But  if  the 
solid  is  made  into  dust  or  the  liquid  is  transformed  into  fine 
droplets,  the  surface  layer  is  enormously  increased  and  new 
ohenomena  become  prominent.     Dusts,  mists,  and  other  finely 
divided  materials  may  be  characterized  as  disperse  systems. 
They  are  to  be  the  subject  of  our  next  study. 

726.  Cohesion  and  Adhesion. — The  structure  of  gases  and 
solids  need  not  be  discussed  further  here.     But  some  points  about 
the  structure  and  behavior  of  liquids  may  well  be  taken  up  in 
preparation  for  the  subsequent  discussion.     We  are  familiar 
with  the  sight  of  a  liquid  falling  in  drops  from  the  tip  of  a  pipette 
or  burette.    As  soon  as  a  given  volume  of  liquid  is  free  from 
outside  influence,  such  as  gravity,  attraction  from  the  walls 
of  the  containing  vessel,  etc.,  it  tends  to  form  a  sphere.     If  the 
drop  is  very  small,  the  sphere  will  be  nearly  perfect  even  when 
pulled  down  against  a  supporting  surface  by  gravity.    This  is 
well  illustrated  in  the  case  of  small  drops  of  mercury  resting  on 
paper.     The  attraction  between  molecules  falls  off  very  rapidly 
as  the  distance  between  them  increases.    As  a  consequence 
only  the  immediate  neighbors  of  a  given  molecule  in  a  liquid 
exert  any  appreciable  influence  upon  it,  and,  except  in  the  surface 
layer,  this  attraction  is  balanced,  since  it  comes  equally  from 
all  sides.     But  in  the  surface  layer  the  molecules  are  pulled 
inward  and  there  is  no  compensating  pull  outward  from  beyond 
the  surface.    The  net  result  is  that  the  surface  layer  acts  as  a 
compressing  membrane  upon  the  rest  of  the  drop.     Work  must 
be  done  against  this  inward  pressure  if  a  liquid  is  to  be  sub- 
divided.   This  fact  may  be  easily  observed  if  the  finger  is  used 
to  divide  a  large  drop  of  mercury  resting  on  a  flat  glass  surface. 

470 


Disperse  Systems  471 

The  floating  of  a  needle  on  the  surface  of  water  is  another  well- 
known  experiment,  used  to  illustrate  this  resistance  of  a  liquid 
to  deformation.  .  When  two  drops  touch  each  other  they 
promptly  merge  into  a  single  drop.  This  occurs  because  at 
the  instant  of  contact  the  surface  layer  no  longer  exists  where  the 
drops  touch.  The  tension  on  the  remaining  surface  layer  forces 
the  drops  together  into  a  new  sphere,  since  the  latter  has  the 
smallest  surface  of  any  shape  which  a  given  volume  of  substance 
may  take. 

The  attraction  of  like  particles  for  each  other  is  called  cohe- 
sion, in  contrast  to  adhesion,  the  attraction  between  unlike 
particles. 

If  a  clean  glass  plate  (free  from  grease)  is  allowed  to  touch 
the  surface  of  water  and  is  then  gently  raised,  the  water  will  be 
lifted  with  it  several  millimeters  before  the  surface  breaks  and 
the  glass  is  released.    The  water  is  said  to  adhere  to  the  glass. 
As  a  matter  of  fact  the  water  molecules  hold  tighter 
to  the  glass  than  to  each  other,  for  when  the  separa- 
tion comes  water  parts  from  water  and  not  from 
glass.    The  surface  of  water  in  a  small,  clean  glass 
vessel  is  concave  upward  because  of  the  spreading  of 
the  water  upon  the  surface  of  the  latter.     If  a  glass 
capillary  tube  is  placed  in  water  the  latter  rises  in 
the  tube  high  above  its  level  outside  (Fig.   103,4).      FIG.  103,4 
The  rising  of  the  water  in  the  tube  against  the  force 
of  gravity  occurs  because  of  the  attraction  of  the  glass  for  the 
water  molecules,  and  stops  when  this  force  is  just  balanced  by 
the  weight  of  water. 

On  the  same  principle,  when  blotting  paper  or  filter  paper  is 
put  into  water  the  latter  rises  in  the  paper  even  against  the  force 
of  gravity.  Before  blotting  paper  was  introduced,  ink  was 
dried  by  dusting  sand  over  it.  The  excess  ink  was  taken  up 
(adsorbed)  on  the  surface  of  the  sand  grains  and  was  removed 
with  the  latter.  The  adsorption  of  liquids  by  fine  powders  has 
become  of  great  importance  in  industry.  Nitrogly  :erine  is 
adsorbed  by  a  very  finely  divided  silica  (silicon  dioxide)  called 
infusorial  earth,  an  important  substance  which  will  be  spoken 


472  Introduction  to  General  Chemistry 

of  again   (732).     In  this  form,  called  dynamite   (692),  nitro- 
glycerine is  much  more  safely  handled  than  as  a  pure  liquid. 

All  liquids  do  not  wet  (adhere  to)  all  solids.     For  example, 
mercury  does  not  wet  glass.     When  mercury  is  contained  in  a 
glass  vessel  the  shape  of  its  surface  is  convex  upward 
(Fig.  1035).      This  curvature  is  due  to  the  usual 
pulling  inward  of  the  molecules  of  the  surface  layer 
of  the  liquid.     If  a  glass  capillary  tube  is  thrust 
into  mercury,  the  level  of  the  latter  inside  the  tube 
is  lower  than  outside  simply  because  the  pull  inward 
of  the  surface  layer  resists  the  deformation  made 
FIG.  1035      by  the  tube.     Mercury  is  not  adsorbed  by  blotting 
paper.     Water,  of  course,  behaves  the  same  way 
toward  surfaces  which  it  does  not  wet.    Apparently  the  force 
of  adhesion  depends  on  the  nature  of  the  substances  in  question. 
727.  Surface  Areas. — The  foregoing  examples  are  sufficient 
to  illustrate  well-known  surface  phenomena.    Table  XXXVI, 
which  gives  the  surface  area  of  a  cube  with  continuous  subdivi- 
sion, illustrates  how  greatly  the  surface  of  a  given  mass  of  sub- 
stance may  vary,   and   consequently  how  effects   which   are 
ordinarily  slight  may  become  of  great  importance. 

TABLE  XXXVI 


Length  of  Edge  of  Cube 

Number  of 
Cubes 

Total  Surface 

i  cm  

I 

6  sq.  cm. 

i  mm  

IC»3 

60  sq.  cm. 

o  .  ooi  mm  

10" 

6  sq.  m. 

0.000,001  mm  

I021 

6000  sq.  m. 

If  these  areas  are  changed  into  the  system  to  which  we  are 
accustomed,  we  find  that  if  a  cube  one  centimeter  on  one  side 
is  divided  into  little  cubes,  each  o.ooi  mm.  on  one  side,  the 
total  surface  is  about  65  square  feet.  If,  however,  the  original 
cube  is  subdivided  into  smaller  cubes,  each  0.000,001  mm.  on 
a  side,  the  total  surface  is  almost  an  acre  and  a  half ! 

728.  Adsorption  of  Gases  by  Charcoal. — If  charcoal  is 
freshly  heated  and  thrust  into  ammonia  gas  confined  over 


Disperse  Systems  473 

mercury,  the  mercury  will  be  seen  to  rise  in  the  tube  as  the  char- 
coal cools  and  finally  to  remain  high  above  its  old  level  (Fig.  104), 
showing  that  the  ammonia  has  been  taken  up  by  the  charcoal. 

There  remains  in  the  latter  the 
cellular  structure  of  the  sub- 
stance from  which  it  was  made, 
so  that  it  is  composed  of  innu- 
merable tiny  pores  and  therefore 
presents  an  enormous  surface. 
The  fast-moving  ammonia  mole- 
cules encounter  this  and  are 

held  upon  it.     A   few  of  the 
FIG.  104 

molecules  are  able  to  escape  the 

carbon,  apparently,  for  all  are  not  taken  up,  and  finally  there  is 
an  equilibrium  between  the  free  ammonia  and  that  adhering 
to  the  charcoal.  The  ammonia  is  said  to  be  adsorbed  by  the 
charcoal.  The  higher  the  temperature,  the  more  rapidly  is 
equilibrium  reached,  but  the  smaller  is  the  amount  adsorbed. 
It  was  on  this  account  that  the  charcoal  was  freshly  heated 
before  it  was  used  in  the  foregoing  experiment,  in  order  that  its 
surface  should  be  free  from  other  gases.  Charcoal  is  a  very 
good  adsorbent  for  gases,  but  it  does  not  adsorb  all  gases  to  an 
equal  degree.  Thus  a  given  sample  was  found  to  adsorb  90 
times  its  own  volume  of  ammonia,  35  times  its  volume  of  carbon 
dioxide,  but  only  i .  7  times  its  volume  of  hydrogen.  In  the 
recent  war,  filters  of  specially  prepared  cocoanut-shell  charcoal, 
mixed  with  other  chemicals,  were  used  in  gas  masks  to  remove 
poison  gases  from  air  inhaled  through  them. 

729.  Air  and  Glass. — So  strongly  is  air  adsorbed  on  the  sur- 
face of  glass  that  great  difficulty  is  found  in  preparing  barometer 
tubes  which  shall  have  a  perfect  vacuum  in  the  space  over  the 
mercury.     It  is  not  sufficient  to  fill  a  glass  tube  with  mercury  and 
then  invert  it  with  the  open  end  under  mercury.    The  tube  is 
filled,  put  under  vacuum,  and  heated  for  some  time  before  the 
air  is  removed  and  the  tube  is  ready  to  be  put  in  place. 

730.  Water  Vapor  and  Glass. — Accurate  workers  have  long 
known  that  glass  surfaces  are  covered  with  a  thin  water  film 


474  Introduction  to  General  Chemistry 

which  is  with  difficulty  removed.  Water  molecules  encounter 
the  glass  surface  and  are  held  there  owing  to  the  very  strong 
attraction  of  water  to  glass,  to  which  attention  has  already  been 
called. 

These  and  other  experiences  show  that  on  the  surface  of  every 
solid  exposed  to  a  gas  we  may  expect  to  find  an  adsorbed  layer 
of  the  latter.  If  the  solids  are  finely  divided,  so  that  the  surface 
is  large,  the  result  may  become  of  great  importance. 

731.  Adsorption  and  Catalysis. — We  have  many  times  noticed 
the  catalytic  effect  of  finely  divided  metals  on  gas  reactions. 
For  example,  the  union  of  hydrogen  and  oxygen  (303)  is  greatly 
accelerated  by  the  presence  of  finely  divided  platinum.    We 
find  that  platinum  takes  up  both  hydrogen  and  oxygen  in  con- 
siderable quantity.     As  early  as  1844  Faraday  pointed  out  that 
one  result  of  the  presence  of  a  metallic  catalyst  is  that  on  its 
surface  the  gases  undergoing  reaction  are  at  far  greater  con- 
centration than  in  the  gas  mixture  itself.     Such  a  layer  would 
react  far  more  rapidly  than  the  main  mixture.    Undoubtedly 
adsorption  is  the  first  stage  of  the  reaction  on  these  contact 
agents.    But  some  absorption,  or  penetration  of  the  adsorbed 
gas  beneath  the  surface  on  which  it  is  at  first  held,  follows,  and 
in  some  cases  compounds  are  formed. 

732.  Adsorption  from  Solution. — If  a  little  of  the  dyes  tuff 
methyl  violet,  often  used  as  the  purple  tint  of  indelible  ink 
pencils,  is  added  to  water,  a  beautiful  purple  solution  results 
which  may  be  filtered  unchanged.     But  if  charcoal  (from  sugar, 
for  instance)  is  shaken  in  this  solution  and  the  mixture  is  then 
poured  on  the  filter,  the  filtrate  is  found  to  be  colorless.    The 
methyl  violet  has  been  adsorbed  by  the  charcoal.    That  the 
dye  has  not  been  destroyed  may  be  shown  by  pouring  alcohol 
through  the  filter.    The  filtrate  again  shows  the  brilliant  color 
of  the  dye.    An  enormous  number  of  substances  are  adsorbed 
from  water  by  charcoal.     It  is  because  of  this  property  that  it  is 
an  effective  filter  for  purifying  water.    However,  there  are  limits 
to  the  amount  that  a  given  column  of  charcoal  can  adsorb,  so 
that  it  soon  loses  this  power.    Bone  black,  charcoal  made  from 
bones,  is  extensively  used  to  take  out  objectionable  impurities 


Disperse  Systems  475 

from  sugar  solution  in  the  process  of  refining  sugar.  Fusel  oil, 
a  poisonous  by-product  present  in  crude  whiskey,  is  also  removed 
from  the  latter  by  filtration  through  charcoal. 

Besides  charcoal  there  are  other  good  adsorbents.  Fullers 
earth,  a  very  fine  clay  (mainly  aluminum  silicate)  which  varies 
in  composition,  is  an  important  industrial  adsorbent  for  the 
purification  of  edible  oils.  The  particles  of  this  earth  are  usually 
between  0.007  of  a  millimeter  and  0.0002  of  a  millimeter  in 
diameter.  Another  important  adsorbent  is  infusorial  earth  or 
kieselguhr,  a  deposit  of  the  skeletons  of  diatoms,  which  are  tiny 
aquatic  organisms.  Beds  of  this  substance  of  as  much  as  a 
thousand  feet  in  thickness  are  found  in  the  United  States.  Other 
adsorbents  are  finely  divided  metals,  plant  and  animal  fibers 
such  as  cotton,  silk,  wool,  etc.  As  has  been  pointed  out  in  the 
case  of  metals  (731),  absorption  may  of  course  follow  adsorp- 
tion in  the  action  of  these  substances.  The  fundamental  require- 
ment for  a  good  adsorbent  seems  to  be  an  enormous  surface  of 
contact  with  the  solution. 

733.  Suspensions   Produced  by  Grinding  of  Solids  under 
Liquids. — If  clay  is  stirred  up  in  water  a  turbid  mixture  results. 
First  the  coarser  particles  settle  to  the  bottom,  and  then  gradually 
finer  and  finer  particles  follow  as  time  goes  on.    Pulverized 
emery  (174),  used  for  grinding,  is  graded  according  to  the  num- 
ber of  minutes  required  for  it  to  settle  after  being  stirred  in 
water.    That  which  will  settle  in  one-half  minute  is  a  coarse 
grade.    Ten-minute  emery  is  for  very  delicate  work.    Careful 
grinding  of  any  substance  will  produce  powders  which  are  still 
slower  in  settling  than  any  of  these,  but  we  have  other  ways  of 
producing  them.  . 

734.  Arsenious  Sulfide  Suspension. — Arsenious  sulfide,  As2S3, 
is  very  insoluble  in  water.    If  this  solid  is  shaken  up  in  pure 
water  only  a  minute  trace  will  be  found  in  a  liter  of  the  latter 
after  filtration.    When  arsenious  acid,  H3As03,  mixed  with  a 
little  hydrochloric  acid,  is  treated  with  hydrogen  sulfide,  a  pre- 
cipitate of  arsenious  sulfide  appears.    But  if  the  hydrochloric  acid 
is  omitted,  only  a  yellow,  opalescent  liquid  results.    When  the 
latter  is  poured  through  filter  paper,  merely  a  trace  of  precipitate 


476  Introduction  to  General  Chemistry 

is  held  back  and  the  liquid  passes  through  unchanged.  That 
the  latter  is  not  a  supersaturated  solution  may  be  easily  estab- 
lished by  adding  to  it  a  little  of  the  solid  arsenious  sulfide.  No 
settling  out  of  the  arsenious  sulfide  follows,  as  would  be  the 
case  if  the  solution  were  supersaturated  (123). 

In  whatever  form  the  arsenious  sulfide  exists  in  the  yellow 
liquid,  its  presence  in  water  seems  to  have  only  a  slight  effect 
upon  the  boiling-point  and  freezing-point  of  the  latter.  The 
osmotic  pressure  (711)  of  such  a  solution  is  very  small.  If  we 
calculate  from  the  latter  the  amount  of  arsenious  sulfide  neces- 
sary to  give  one  gram  molecular  weight,  we  obtain  enormous 
numbers;  for  instance,  in  one  case  six  thousand  grams  was  the 
result,  a  number  which  is  more  than  twenty-four  times  that 
indicated  by  the  formula  As2S3. 

The  examination  of  the  liquid  under  the  ordinary  microscope 
shows  nothing.  But  if  it  is  examined  even  with  a  crude  form 
of  the  ultra-microscope  described  in  706,  it  is  found  to  be  full  of 
dancing  particles  which  must  be  arsenious  sulfide.  All  the 
evidence  shows  that  the  arsenious  sulfide  exists  in  the  form  of 
particles  which  are  very  large  compared  to  the  simple  arsenious 
sulfide  molecules,  but  small  compared  to  particles  which  settle 
from  solution. 

735.  Colloids. — Such  a  non-settling  suspension  is  referred  to 
as  a  colloidal  solution,  or  as  a  colloidal  suspension.  Usually 
matter  is  said  to  be  in  a  colloidal  state  if  it  is  too  finely  divided 
to  be  held  back  by  a  good  filter  paper  and  still  coarse  enough 
to  be  seen  in  the  ultra-microscope.  This  means  that  the  average 
diameter  of  the  particles  is  between  o  .000,1  and  o  .000,001  mm. 
The  setting  aside  of  these  systems  as  a  special  class  is  obviously 
entirely  a  matter  of  convenience.  Colloidal  suspensions  of 
substances  which  are  solids  when  in  massive  form  are  often  called 
suspensoids,  while  colloidal  suspensions  of  liquids  are  called 
emulsoids.  The  name  colloid,  which  means  glue-like,  was 
originated  by  Thomas  Graham,  the  first  important  investigator 
of  this  subject.  He  worked  mainly  with  gum  arabic,  starch, 
glue,  and  glue-like  substances  which  belong  to  a  more  complex 
type  than  we  have  yet  studied. 


Disperse  Systems 


477 


736.  Properties  of  Colloidal  Arsenious  Sulfide,  Diffusion. — 

The  most  striking  thing  about  the  colloidal  arsenious  sulfide 
suspension  is  its  stability.  If  it  is  kept  in  good-grade  glass 
bottles,  months  may  pass  without  its  settling. 

If  two  test  tubes  are  half  filled  with  5  per  cent  hot  gelatin 
solution  which  is  allowed  to  cool  and  set  to  a  gel,  and  to  one  tube 
is  added  a  layer  of  copper  sulfate  solution,  while  to  the  other  the 
yellow  arsenious  sulfide  suspension  is  added,  after  about  twenty- 
four  hours  the  blue  solution  will  have  penetrated  well  into  the 
gel  layer,  but  the  yellow  solution  will  not  have  entered  the  layer 
below  it.  The  effect  can  best  be  seen  by  corking  the  tubes  and 
then  inverting  them ,  The  gel  in  the  copper  sulfate  tube  will  be 
found  to  be  partly  colored,  while  the  other  gelatin  layer  will 
be  found  to  be  unchanged.  Apparently  the  rate  of  diffusion 
of  the  arsenious  sulfide  particles  is  very  slow.  This  is  explained 
by  the  work  of  Perrin  (707).  Each  arsenious  sulfide  particle 
is  very  much  larger  than  a  copper  sulfate  molecule  or  a  copper 
ion,  and  since  it  is  at  the  same  temperature  its  velocity  must  be 
much  smaller. 

737.  Effect  of  an  Electrical  Current  on  Colloidal  Arsenious 
Sulfide. — Colloidal  arsenious  sulfide  solution  is   a  very  poor 
conductor  of  electricity,  but  if  we  fill  a  U-tube 

with  this  material  and  pass  a  no- volt  current 
through  it,  in  from  ten  to  twenty  minutes  a 
colorless  layer  will  be  plainly  visible  at  the 
negative  electrode,  and  the  region  near  the 
positive  electrode  will  be  deepened  in  color. 

The  effect  is  better  shown  by  using  the 
device  described  by  Professor  A.  A.  Noyes 
and  shown  in  Fig.   105.     The  ends  of  the 
inner  tube  are  covered  with  thin  parchment 
paper,  or  better  with  goldbeater's  skin.     Over 
the  ends  are  fitted  extension  tubes,  which  are 
joined  to  the  U-tube  bv  rubber.    The  device  is 
inverted  for  filling  through  the  hole,  which  is  finally  closed  by 
sliding  a  rubber  tube,   H,  over  it.     The  spaces  around   the 
electrodes  are  filled  with  ordinary  distilled  water.     The  shaded 


FIG.  105 


478  Introduction  to  General  Chemistry 

area  of  the  figure  shows  the  region  of  yellow  color  after  the 
current  has  been  passing  for  some  time. 

Since  the  suspension  drifts  away  from  the  negative  and 
toward  the  positive  electrode,  its  particles  are  negatively 
charged.  We  have  already  noticed  that  when  unlike  substances 
are  in  contact,  a  loss  of  electrons  by  one  and  a  gain  of  electrons 
by  the  other  are  very  likely  to  occur  (474).  It  would  not  be 
surprising  if  this  were  the  cause  of  the  existence  of  this  difference 
of  potential  between  the  solid  and  the  liquid.  But  we  shall 
soon  see  that  other  factors  may  also  account  for  this  phenomenon. 

738.  Effect  of  Electrolytes  upon  Colloidal  Arsenious  Sul- 
fide. — We  have  already  seen  that  in  the  presence  of  hydrochloric 
acid  the  colloid  does  not  form,  and  we  find  that  if  a  little  hydro- 
chloric acid  is  added  to  the  suspension  the  liquid  becomes  very 
much  more  turbid,  and  with  further  addition  of  the  acid  a  pre- 
cipitate soon  appears  and  settles  to  the  bottom  of  the  container. 
The  curious  thing  is  that  any  other  good  electrolyte  will  accom- 
plish the  same  thing. 

The  positive  ion  seems  to  be  the  active  part  of  the  electrolyte; 
for  about  the  same  concentrations  of  salts  which  have  univalent 
positive  ions  are  necessary  to  cause  complete  coagulation  of  the 
precipitate  from  a  given  volume  of  solution.  Smaller  concen- 
trations of  salts  which  have  bivalent  positive  ions  are  needed, 
and  very  much  smaller  concentrations  of  salts  which  have 
trivalent  positive  ions.  Relatively  little  difference  is  made  by 
varying  the  negative  ion  of  the  salt. 

739.  Adsorption  of  the  Precipitating  Agent. — Careful  experi- 
ments have  shown  that  the  active  ion  is  carried  down  by  the 
precipitate,  leaving  in  solution  an  equivalent  amount  of  the 
corresponding  ion  of  water;    thus  arsenious  sulfide  coagulated 
with  calcium  chloride  contained  calcium,  and  the  liquid  left 
behind   contained   an   equivalent   concentration   of   hydrogen 
chloride. 

Apparently  the  positive  ion  neutralizes  the  charge  on  the 
suspension  and  is  carried  down  by  the  precipitate.  Since  the 
proportion  of  ion  carried  down  by  the  precipitate  bears  no  simple 
and  constant  relation  to  the  weight  of  the  precipitate,  as  would 


Disperse  Systems  479 

be  the  case  if  a  chemical  compound  of  the  type  we  know  had 
been  formed,  it  is  apparent  that  the  product  is  of  a  different 
class.  For  the  sake  of  convenience  we  may  call  the  former  an 
adsorption  compound.  If  the  adsorption  compound  is  washed 
with  pure  water,  the  calcium  is  not  removed,  but  it  can  be  taken 
away  by  washing  the  substance  with  some  other  electrolyte, 
ammonium  chloride,  for  example.  In  this  case  ammonium 
takes  the  place  of  calcium  with  the  arsenious  sulfide,  and  the 
calcium  is  found  as  calcium  chloride  in  solution. 

The  proportion  of  the  adsorbed  ion  to  the  precipitate  is 
usually  very  small,  though  of  course  different  with  different  prep- 
arations according  to  the  valence  of  the  ion,  the  charge  on  the 
suspensions,  and  other  conditions.  For  example,  in  the  case  of 
certain  colloidal  arsenious  sulfi.de  suspensions  coagulated  with 
calcium  chloride  the  proportion  was  found  to  be  one  equivalent 
of  calcium  to  fifty  equivalents  of  the  sulfide.  If  very  pure  sub- 
stances are  desired,  as  is  the  case  in  exact  analysis,  even  this 
small  proportion  may  be  of  moment,  especially  if  the  adsorbed 
ion  has  a  high  molecular  weight. 

Experience  shows  that  the  adsorption  of  ions  from  solution 
by  a  precipitate  in  the  processes  of  its  formation  is  the  rule  and 
always  must  be  taken  into  consideration  in  the  preparation  of 
pure  substances. 

740.  Influence  of  the  Charge  on  the  Stability  of  a  Suspen- 
soid. — If  in  the  experiment  on  the  migration  of  the  suspensoid 
(737)  the  particles  of  the  latter  had  been  allowed  to  reach  the 
positive  electrode,  they  would  have  been  precipitated.    Appar- 
ently the  existence  of  the  charge  on  the  suspension  is  necessary 
for  its  stability.    We  can  understand  that  the  presence  of  like 
charges  on  the  arsenious  sulfide  would  have  a  tendency  to  keep 
them  from  coming  together,  but  undoubtedly  the  effect  of  the 
charge  is  more  complex  than  this. 

741.  Preparation  of  Colloidal  Ferric  Hydroxide. — Another 
method  of  preparing  colloidal  solutions  is  illustrated  in  the  follow- 
ing preparation  of  colloidal  ferric  hydroxide.     Ferric  chloride 
solution,  which  is  of  course  acid  in  reaction  (436),  is  treated  with 
ammonium  carbonate  as  long  as  the  precipitate  which  first 


480  Introduction  to  General  Chemistry 

forms  redissolves.  The  dark-red  solution  thus  made  contains 
mainly  ferric  chloride,  colloidal  ferric  hydroxide,  and  ammonium 
chloride.  The  ferric  hydroxide  and  ammonium  chloride  are 
produced  as  the  ammonium  carbonate 
reacts  with  the  hydrochloric  acid  formed 
by  the  hydrolysis  of  ferric  chloride : 

FeCl3+3H2O^Fe(OH)3+3HCl. 

This  mixture  is  next  placed  in  a  parchment- 
paper  bag  (Fig.  106),  and  the  whole  is  sus- 
pended in  a  water  bath  through  which  a 
current  of  fresh  water  is  constantly  flow- 

FlG.    I 06  J 

ing.  If  this  arrangement  is  left  for  about 
four  days  the  contents  of  the  bag  will  be  found  to  give  only 
a  very  small  test  for  chloride,  showing  that  virtually  all  of  the 
ferric  chloride  has  been  transformed  into  ferric  hydroxide. 
The  preparation  is  in  fact  colloidal  ferric  hydroxide. 

742.  Properties  of  Colloidal  Ferric  Hydroxide. — When  the 
liquid  is  poured  through  filter  paper  no  precipitate  is  left  on  the 
latter.  The  freezing-point,  boiling-point,  and  osmotic-pressure 
determinations  all  show  that  the  substance  is  present  in  the 
form  of  particles  which  are  large  in  comparison  with  simple 
molecules.  The  dark-red  liquid  is  stable  if  carefully  kept.  If 
it  is  placed  in  the  apparatus,  as  shown  in  Fig.  104,  and  an  electric 
current  is  applied,  the  red  substance  is  found  to  migrate  toward 
the  negative  electrode,  proving  that  this  suspension  is  positive 
to  its  solution. 

Electrolytes  precipitate  the  red  hydroxide,  leaving  the  liquid 
colorless.  This  time  it  is  the  negative  ion  of  the  electrolyte 
which  is  active  and  goes  to  form  the  adsorption  compound. 
Again  the  valence  of  the  precipitating  ion  is  found  to  be  an 
important  factor.  In  a  given  instance  about  one-fortieth  as 
great  a  concentration  of  sulfate  as  chloride  ion  was  needed  to 
coagulate  a  given  amount  of  ferric  hydroxide. 

When  the  latter  suspension  is  added  a  little  at  a  time  to 
arsenious  sulfide  suspensions,  the  mixture  becomes  turbid,  and 
finally  a  precipitate  of  the  two  substances  settles  out.  This  is 


Disperse  Systems  481 

the  result  of  mutual  adsorption  and  precipitation.  Curiously 
enough  the  addition  of  a  large  excess  of  either  colloid  does  not 
give  a  precipitate,  but  gives  a  complex  colloid. 

743.  Explanation   of   the   Preparation   of   Colloidal   Ferric 
Hydroxide. — If  a  layer  of  the  colloidal  ferric  hydroxide  is  put 
over  a  layer  of  gel,  as  was  described  in  the  case  of  arsenious 
sulnde  (736),  the  ferric  hydroxide  will  be  found  to  diffuse  very 
slowly  indeed.    This  fact  was  made  use  of  in  the  original  prepara- 
tion of  the  colloidal.     The  ferric  hydroxide  made  by  the  action 
of  ammonium  carbonate  did  not  diffuse  through  the  parchment 
paper  with  appreciable  speed,  and  at  the  same  time  it  adsorbed 
the  ferric  chloride  so  that  the  latter  remained  in  the  bag.     The 
hydrochloric  acid  and  the  ammonium  chloride,  however,  passed 
readily  through  the  paper  and  were  washed  away  on  the  other 
side.     The  continuous  loss  of  hydrochloric  acid  allowed  the 
hydrolysis  of  the  ferric  chloride  to  go  to  completion.    The 
separation  of  colloids  from  dissolved  substances  by  a  process 
of  diffusion  through  a  membrane  is  called  dialysis  and  the 
apparatus  is  called  a  dialyzer. 

744.  Colloidal  Silver. — If  an  electric  arc  is  passed  between 
two  silver  wires  submerged  in  water,  a  dark  cloud  will  form 
around  the  electrodes.    This  is,  in  fact,  a  colloidal  suspension 
of  silver.    A  little  alkali  added  to  the  liquid  will  increase  the 
stability  of  the  suspension.    The  dark  liquid  has  the  same  general 
properties  as  those  of  the  other  suspensions  we  have  described. 
It  is  electro-negative  to  water.    This  method  of  producing  a 
colloidal  suspension  of  a  metal  by  volatilizing  the  latter  in  the 
electric  arc  is  named  Bredig's  method  after  its  discoverer,  the 
German  chemist  Bredig. 

745.  Protecting  Agents. — The  silver  suspension  can  be  made 
of  considerable  concentration  if  about  i  per  cent  of  gelatin  is 
added  to  the  solution.    The  effect  of  the  gelatin  may  be  shown 
by  mixing  silver  nitrate  (N/io)  and  hydrochloric  acid  (N/io) 
solutions,  to  each  of  which  about  i  per  cent  of  gelatin  has  been 
freshly  added.    Instead  of  the  copious  white  precipitate  which 
we  have  so  often  seen  result  from  mixing  these  reagents,  only  a 
slight  white  turbity  appears.    The  gelatin  has  prevented  large 


482  Introduction  to  General  Chemistry 

particles  from  forming.  Many  other  agents  besides  gelatin 
stabilize  suspensions.  Their  function  is  not  understood,  but  it  is 
thought  that  they  surround  the  particles,  thus  preventing  further 
union  between  them.  Hence  they  are  called  protecting  agents. 

746.  Red   Gold   Suspensions. — A   very   useful   reagent  in 
making  suspensions  of  metals  is  tannin,  a  complex  organic  sub- 
stance which  is  both  a  good  reducing  agent  and  a  protecting 
agent.     A  very  finely  divided  gold  suspension  may  be  made 
with  a  neutralized  (about  i  per  cent)  gold  chloride  solution  and 
dilute  (o .  i  per  cent)  solution  of  tannin,  according  to  a  method 
described  by  Dr.  Wolfgang  Ostwald.     First  a  few  drops  of  the 
gold  solution  are  mixed  with  100  c.c.  of  water,  then  a  few  drops 
of  the  tannin  solution  are  added,  and  the  mixture  is  heated  for 
a  few  minutes,  with  constant  shaking.     Meantime  the  red  celor 
of  the  gold  suspensoid  appears.     More  gold  chloride  and  tannin 
may  be  added  alternately  a  little  at  a  time. 

The  particles  of  the  red  gold  suspension  are  very  tiny,  usually 
about  one  or  two  one-hundred-thousandths  of  a  millimeter  in 
diameter.  These  were  the  particles  described  by  Zsigmondy 
(706).  They  are  negative  to  the  water  in  which  they  are 
suspended  and  are  precipitated  by  positively  charged  ions. 
Coarser  suspensions  of  gold  may  be  violet  or  blue  in  color. 
Brown  gold  suspensions  settle  in  a  very  short  time. 

747.  Summary  of  Work  with  Suspensoids. — The  general 
methods  of  preparing  suspensoids  are  illustrated  in  the  foregoing. 
Either  the  material  is  subdivided  by  grinding  or  volatilization 
in  the  electric  arc,  or  it  is  formed  in  solution  and  the  particles 
are  not  permitted  to  grow  to  a  size  large  enough  to  settle  from 
the  liquid.    The  absence  of  electrolytes  and  the  presence  of  pro- 
tecting  agents   assist    in    making    concentrated    suspensions. 
Colloidal  suspensions  of  an  enormous  number  of  different  sub- 
stances have  been  prepared.    Apparently  any  substance  can 
exist  in  the  colloidal  state  in  liquids  in  which  it  is  not  soluble. 
Thus  colloidal  sodium  suspension  has  been  prepared  in  ether, 
phosphorus  in  water,  sodium  chloride  in  benzene,  etc. 

Apparently  many  inorganic  substances  stabilize  colloidal 
suspensions  to  a  moderate  degree,  for  we  often  see  opalescent 


Disperse  Systems  483 

liquids  form  when  precipitation  in  dilute  solutions  is  in  progress. 
As  the  last  of  the  precipitating  agent  is  added,  the  solution  be- 
comes clear  and  the  precipitate  settles.  Thus  if  5  c.c.  of  N/2O 
silver  nitrate  is  added  to  20  c.c.  of  N/io  potassium  chloride  solu- 
tion, a  part  of  the  silver  chloride  appears  as  a  non-settling 
suspension.  If  more  and  more  silver  nitrate  is  added,  with 
constant  shaking  of  the  mixture,  the  liquid  finally  becomes 
transparent  and  the  precipitate  settles  just  as  the  amount  of 
silver  nitrate  added  becomes  nearly  equivalent  to  the  potassium 
chloride.  The  excess  potassium  chloride  present  at  first  stabi- 
lizes the  silver  chloride  suspension. 

The  origin  of  the  charges  on  these  suspensions  is  not  under- 
stood. As  has  been  said,  it  would  not  be  surprising  to  find  a 
frictional  charge  on  them,  but  we  must  also  consider  that  an 
unequal  adsorption  of  positive  and  negative  ions  from  solution 
might  also  account  for  the  existence  of  these  charges.  In  addi- 
tion, the  loss  of  ions  from  the  suspended  particles  to  the  solution 
might  occur.  The  fact  that  all  suspensions  of  bases  are  positive 
to  the  liquid  in  which  they  exist  would  seem  to  favor  the  idea 
of  the  ionization  of  the  solid  particles  in  these  cases.  Thus  the 
hydroxide  particle  might  lose  one  or  more  hydroxyl  ions  to  the 
solution,  leaving  a  residue  which  would  be  a  very  large  positive 
ion.  But  in  explaining  the  presence  of  these  charges  we  must 
also  consider  that  most  other  suspensions  in  water  are  negative. 
Thus,  finally,  suspended  clay,  lamp  black,  metals,  sulfur,  salts, 
etc.,  are  negative  to  their  solutions  as  a  rule,  though  not  always. 
Very  probably  all  three  possible  causes  cited  for  the  existence 
of  the  charges  come  into  play  at  different  times. 

748.  Tyndall  Effect. — When  a  pencil  of  light  from  a  lantern 
is  brought  to  bear  on  water,  we  see  only  a  faint  glow  over  its 
path  through  the  latter;  but  as  soon  as  a  colloidal  suspension 
is  added,  the  path  of  the  light  ray  appears  as  a  cone  of  bright 
cloud.  This  effect  is  called  the  Tyndall  effect  in  honor  of 
J.  Tyndall,  who  was  the  first  to  make  extensive  use  of  this 
phenomenon.  The  Tyndall  effect  is,  of  course,  best  observed 
in  a  darkened  room.  It  depends  on  the  difference  in  refracting 
power  for  light  of  the  particles  of  the  suspension  and  of  the  pure 


484  Introduction  to  General  Chemistry 

liquid.  In  some  cases  this  difference  is  not  great,  so  that  no 
cone  appears.  Hence  the  absence  of  the  Tyndall  cone  does  not 
always  mean  that  the  liquid  in  question  contains  no  colloidal 
matter.  The  ultra-microscope  is  of  course  simply  a  refined 
apparatus  for  examining  the  Tyndall  effect  in  detail.  It  is 
interesting  to  note  that  concentrated  solutions  of  sugar,  sodium 
acetate,  and  many  other  salts  show  the  Tyndall  effect. 

749.  Test  for  the  Charge  on  a  Suspension. — A  crude  distinc- 
tion may  be  made  between  positive  and  negative  colloids  as 
follows:  Strips  of  filter  paper  are  suspended  so  that  one  end  of 
each  strip  is  wet  by  a  colloid.     If  a  strip  dips  into  an  arsenious 
sulfide  or  gold  solution,  the  liquid  which  rises  in  the  paper  will 
show  the  color  of  the  colloid.    The  colloid  is  spreading  through 
the  paper,  though  usually  not  so  rapidly  as  the  water.    But  if 
the  strip  dips  into  a  ferric  hydroxide  solution,  the  ferric  hydroxide 
will  be  found  to  diffuse  but  a  little  way  into  the  paper,  although 
water  from  the  suspension  rises  through  the  latter  as  easily  as 
in  the  other  cases.    The  test  is  reliable  only  under  average  con- 
ditions, which  are  those  of  dilute  negative  suspensions  and 
relatively  more  concentrated  positive  suspensions.    If  turbid 
suspensions  are  to  be  tested,  they  should  be  filtered  before  the 
trial  is  made,  otherwise  coarse  particles  will  impede  the  rise  of 
the  colloid  through  the  paper. 

750.  Important  Suspensoids. — Suspensoids  have  become  of 
considerable    practical    importance.     Colloidal    silver    is    an 
important  antiseptic.    It  is  prepared  with  a  protecting  agent 
so  that  it  may  be  sold  as  a  dry  powder  (argyrol),  which  forms  a 
dark-brown  colloid  upon  the  addition  of  water.     Colloidal  copper, 
mercury,  and  sulfur  have  also  come  into  use  in  medicine.     India 
ink  is  mainly  colloidal  carbon.   The  important  lubricants  aquadag 
and  oildag  are  colloidal  graphite  in  water  and  oil,  respectively, 
each  protected  by  tannin.    The  lubricating  power  of  graphite  is 
much  improved  when  it  is  in  the  colloidal  state. 

We  shall  next  take  up  emulsoids.    These,  as  we  shall  see,  are 
of  tremendous  importance  in  biochemistry. 

751.  Oil  and  Water. — If  oil,  benzene  for  example,  is  shaken 
up  in  water,  for  a  few  seconds  there  is  a  general  mixture  of  oil 


Disperse  Systems  485 

and  water.  Then  two  cloudy  layers  separate.  The  lower  is  an 
emulsion  of  benzene  in  water,  and  the  upper  an  emulsion  of 
water  in  benzene.  Very  soon,  however,  the  two  layers  become 
transparent.  Momentarily  the  oil  and  water  are  in  the  colloidal 
state,  but  this  is  not  permanent.  If  a  drop  of  water  touches 
another  drop  of  water,  the  two  promptly  coalesce  (726).  The 
oil  has  very  little  attraction  for  the  water,  and  its  drops  continue 
uniting  just  as  the  water  drops  do  until  two  layers  are  made  from 
all  the  small  drops.  If,  however,  soap  (678)  is  added  and  the  oil 
and  water  layers  are  shaken  together,  a  stable  suspension  of  oil 
in  water  will  form.  The  suspension  may  be  termed  an  emulsoid 
if  the  droplets  are  small  enough  (735). 

752.  Cleansing  Action  of  Soap. — Apparently  the  emulsifying 
power  of  soap  is  an  important  factor  in  its  cleansing  or  detergent 
action.    Thus  a  little  fine  dirt  shaken  with  water  settles  out, 
but  shaken  with  soap  and  water  it  remains  in  suspension.     Dirt 
usually  sticks  to  the  soap  solution  more  than  it  does  to  the 
fabric  being  washed,  and  hence  rinsing  carries  off  the  dirt. 
Undoubtedly  some  hydrolysis  of  soap  in  water  does  occur,  since 
soaps  are  all  salts  of  the  very  weak  fatty  acids  (678),  but  the 
alkali  formed  is  not  an  important  factor  in  the  detergent  action 
of  soap,  since  alkali  alone  has  no  such  power  to  emulsify  mineral 
oils,  etc.     Vegetable  oils  can  be  emulsified  to  some  extent  with 
dilute  alkalies.     But  it  must  be  remembered  that  these  are 
esters  (677)  and  that  some  soap  is  formed  by  the  interaction  of 
esters  with  the  alkali.     If  alkali,  washing  soda,  or  ammonia  is 
added  to  soap  solution,  its  detergent  power  is  increased.     Of 
course  the  hydrolysis  of  soap  is  decreased  at  the  same  time. 

753.  Nature  of  Soap  Solution. — Ultra-microscopic  examina- 
tion of  soap  solution  is  unsatisfactory,  presumably  because  there 
is  very  little  difference  in  the  power  of  soap  and  water  to  refract 
light.     The  fact  that  soap  has  little  effect  on  the  boiling-point 
of  water  seems  to  indicate  a  very  high  molecular  weight.     All 
evidence  points  to  the  conclusion  that  the  soap  molecules  in 
solution  are  groups  of  large  numbers  of  simple  molecules. 

754.  Work  of  Harkins  and  of  Langmuir. — It  has  long  been 
known  that  a  soap  solution  is  more  concentrated  in  the  surface 


486  Introduction  to  General  Chemistry 

layer  than  in  the  inner  region.  Professor  W.  D.  Harkins  and 
co-workers  have  found  that  this  is  true  of  solutions  of  sodium 
oleate  (678,  681)  as  dilute  as  0.002  normal.  The  surface  layer 
of  a  0.002  normal  solution  is  saturated.  Stronger  solutions 
up  to  o.i  normal  have  the  same  concentration  in  the  surface 
layer  as  have  these  very  dilute  solutions.  In  each  case  there  is 
equilibrium  between  molecules  of  soap  in  the  surface  layer  and 
in  the  body  of  the  solution.  Evidently  the  surface  layer  allows 
the  escape  of  relatively  few  molecules  to  the  solution,  or  the 
inequality  of  concentration  could  not  be  kept  up.  Further, 
Professor  Harkins  and  Dr.  Langmuir,  of  the  General  Electric 
Company,  working  independently  and  by  different  methods, 
have  shown  that  the  soap  molecules  of  the  surface  layer  are  not 
simply  jumbled  together  but  are  arranged  in  a  definite  order, 
The  long  sodium  oleate  molecules 

OHHHHHHHHHHHHHHHHH 

II!  Ill  I        I    I    I    II 

Na-O-C-C-C-C-C-C-C-C-C  =  C-C-C-C-C-C-C-C-C-H 

I     I     I     I     I  I  III  I 

HHHHHHH  HHHHHHHH 

O 

II 
are  all  placed  with  the  (Na  •  0  •  C  • )  group  toward  the  body  of  the 

solution  and  the  long  hydrocarbon  chain  (664)  outward. 

It  is  an  old  and  well-known  rule  that  "like  dissolves  like." 
Thus  oils  are  good  solvents  for  fats  and  greases  but  poor  solvents 
for  sugar,  while  water  is  a  good  solvent  for  the  latter  but  a  poor 

O 

solvent  for  the  former.  Apparently  water  attracts  the  Na  •  O  •  C 
group  of  the  soap  molecule  strongly  but  has  no  attraction  for  the 
long  hydrocarbon  chain.  As  a  result  of  this  attraction  on  one 
end  of  the  long  molecule,  this  end  is  pulled  in  toward  the  liquid, 
and  the  hydrocarbon  chain  is  left  on  the  surface.  The  surface 
of  the  soap  solution  is  therefore  that  of  an  oil,  and  it  is  on  this 
account  that  soap  solutions  can  wet  ojls  while  water  cannot. 

755.  Other  Emulsifying  Agents. — Other  emulsifying  agents 
are  in  common  use,  but  their  function  is  not  so  well  understood. 


Disperse  Systems  487 

Thus  in  mayonnaise  dressing  the  yolk  of  an  egg  is  used  to  keep 
the  oil  and  vinegar  together.  Many  substances  are  capable  of 
forming  emulsoids  without  the  assistance  of  special  agents. 
Examples  of  these  are  gelatin,  gum  arabic,  egg  white,  etc.  These 
are  all  substances  which  are  themselves  wet  by  water. 

756.  General    Properties    of    Emulsoids. — Emulsoids    are 
sharply  differentiated  from  suspensoids  by  their  relatively  high 
viscosity.     The  tiny  particles  of  both  show  Brownian  move- 
ments (705).     Unlike  suspensoids,  many  emulsoids  seem  to  be 
without  electrical  charges.     Salts  in  small  amounts  have  little 
effect  on  these  systems.     In  large  amounts  they  may  cause 
precipitation.     Thus  soaps  are  precipitated  ("salted  out")  from 
solution  upon  the  addition  of  considerable  quantities  of  sodium 
chloride.     Although  we  have  discussed  emulsions  in  water  only, 
it  is  obvious  that  they  may  exist  in  other  media. 

757.  Important  Emulsoids. — Virtually  all  fluids  of  plant  and 
animal  bodies  are  emulsoids.     Thus  the  sap  of  plants  and  the 
blood  and  milk  of  animals  are  complex  emulsoids. 

758.  Gels. — If  the  solvent  is  evaporated  from  a  suspensoid, 
the  latter  is  left  as  a  powder.    But  if  the  solvent  is  evaporated 
from  an  emulsoid,  or  if  a  concentrated  emulsoid  is  cooled,  a  gel 
usually  forms,  though  not  always.     Soap  melted  in  hot  water 
sets  as  a  gel  on  cooling.     Solutions  of  gelatin  of  greater  concen- 
tration than  0.25  per  cent  will  set  at  temperatures  above  o°. 
These  gels  are  stiffer  the  higher  the  concentration  of  the  original 
emulsion. 

Unfortunately  they  refract  light  little  differently  from  water, 
so  that  an  examination  of  their  structure  under  the  ultra-micro- 
scope is  not  satisfactory.  Gels  have  been  precipitated  by  the 
addition  of  alcohol,  etc.,  and  then  examined.  Their  structure 
then  seems  to  be  that  of  a  delicate  latticework  or  honeycomb, 
but  it  is  more  likely  that  this  form  is  acquired  in  the  process  of 
precipitation. 

We  have  already  seen  that  substances  in  solution  diffuse 
through  gels  very  readily.  In  most  cases  they  do  so  as  easily 
as  through  water,  but  colloidal  substances,  either  emulsoid  or 
suspensoid,  diffuse  through  gels  only  very  slowly.  Some  gels 


488  Introduction  to  General  Chemistry 

are  elastic  and  others  are  not.  If  the  former  are  stretched  they 
become  warm,  and  cool  again  on  contraction.  Another  very 
striking  property  of  gels  is  their  ability  to  absorb  (731)  water. 
If  a  warm,  concentrated  gelatin  solution  is  poured  out  on  a  glass 
plate  and  allowed  to  set,  it  may  be  cut  into  pieces  of  equal  volume. 
The  latter  should  then  be  dried.  If  a  piece  is  allowed  to  soak 
in  water,  it  will  be  found  to  swell.  If  at  the  same  time  other 
pieces  of  the  same  weight  are  allowed  to  soak  about  twenty-four 
hours  in  dilute  acid  or  dilute  alkali  of  equivalent  concentration, 
the  rate  of  swelling  of  the  gel  will  be  found  to  be  most  in  the  acid 
solution,  next  in  the  alkali,  and  least  in  the  pure  water.  Salts 
added  to  pure  water  may  increase  or  decrease  the  rate  of  swelling. 
Still  another  important  property  of  gels  is  their  separation  into 
two  layers,  one  liquid  and  the  other  gel.  If  evaporation  is 
prevented,  all  gels  act  in  this  way  in  the  course  of  time. 

759.  Plant  and  Animal  Tissue. — Virtually  all  of  the  firmer 
parts  of  animals  and  plants  are  gels  of  great  complexity,  but  the 
properties  given  above  for  gelatin  belong  to  most  of  these  sub- 
stances.    For  example,  the  experiment  with  the  gelatin  squares 
may  be  repeated  with  animal  tissue  (frogs'  legs,  for  example) 
with  the  same  type  of  result.     The  development  of  a  watery 
fluid  by  a  gel  is  repeated  in  the  formation  of  many  animal  and 
plant  secretions. 

760.  Rubber. — If  rubber,  which  is  a  typical  colloid,  is  allowed 
to  soak  in  benzene  or  carbon  disulfide,  it  swells  enormously,  at 

the  same  time  taking  up  the  liquid. 
Upon  being  stretched  rubber  becomes 
warm,  and  cools  when  contracting,  just 
as  do  other  gels.  An  interesting  appli- 
cation of  the  general  law  governing  all 
changes,  applications  of  which  we  have 
observed  so  many  times  (367),  may  be 
shown  in  the  following  experiment:  A 
FIG.  107  stout  rubber  band  is  attached  by  one 

end  to  a  support  and  is  stretched  by  a 

weight,  as  in  Fig.  107.  If  a  lighted  match  is  held  near  the 
rubber,  so  that  it  is  heated  quickly  but  not  melted,  the  weight 


Disperse  Systems  489 

will  be  seen  to  rise.    The  application  of  heat  has  favored  the 
change  which  absorbs  heat,  namely,  contraction. 

761.  Inorganic    Gels. — Many    inorganic    substances    form 
gels  easily;    for  example,   silicic   acid,   ferric  hydroxide,   and 
aluminium  hydroxide.    These  are  all  extremely  insoluble  sub- 
stances.   Von  Weimarn  has  shown  that  in  general,  if  a  very 
high  degree  of  supersaturation  is  attained  preliminary  to  pre- 
cipitation, the  precipitate  will  be  a  gel.    Thus  barium  sulfate 
is  usually  seen  to  precipitate  as  a  powder.     But  if  saturated 
solutions  of  the  very  soluble  salts  sodium  sulfate  and  barium 
sulfocyanate  are  mixed,  a  gel  of  barium  sulfate  forms.     Con- 
centrated solutions  of  sodium  carbonate  and  calcium  chloride 
give  a  gel  of  calcium  carbonate  instead  of  the  usual  powder. 

762.  Relative    Stability    of    Precipitates. — Apparently    the 
crystalline  form  is  more  stable  than  either  the  gel  or  the  powder 
precipitate.    On  long  standing  in  contact  with  the  solution  most 
non-crystalline  precipitates  become  crystalline.    The  rate  of 
change  is  accelerated  by  a  rise  of  temperature  which  increases 
the  rate  of  molecular  agitation  and  usually  also  increases  the 
solubility  of  the  precipitate.     Even  such  substances  as  gelatin 
can  be  prepared  in  crystalline  form.    A  substance  can  separate 
from  solution  in  the  crystalline  form  only  if  there  is  time  and 
opportunity  for  the  orderly  arrangement  of  molecules  in  crystals. 
Hence  only  the  more  soluble  substances  usually  appear  in  this 
form. 

763.  Explanation  of  Adsorption  from  Solution. — The  fact 
that  soap  is  more  concentrated  in  the  surface  layer  than  in  the 
solution  is  a  phenomenon  of  a  type  which  is  very  common  among 
substances  of  complex  molecular  structure.    Most  emulsoids 
and  suspensoids  show  this  to  a  marked  degree,  though  substances 
in  solution  do  also  to  some  extent.     Under  ordinary  conditions 
this  is  not  noticeable.    But  when  the  surface  of  the  solution  in 
question  is  enormously  increased  by  being  mixed  with  a  porous 
or  finely  divided  solid  like  charcoal  or  infusorial  earth,  the  net 
result  is  that  a  great  amount  of  the  colloid  goes  to  the  surface 
layer  and  adheres  to  the  solid  with  the  latter,  while  the  rest  of 
the  liquid  is  drained  away.    The  very  large  extent  of  the  surface 


490  Introduction  to  General  Chemistry 

means  a  very  great  loss  of  material  from  the  original  liquid. 
The  latter  may  be  virtually  freed  from  the  colloidal  matter  if 
passed  through  a  thick  layer  of  filter. 

764.  Importance  of  Colloid  Chemistry. — The  many  complex 
systems  discussed  in  this  chapter  have  one  thing  in  common — 
an  enormous  surface  of  contact  between  different  simpler 
systems.  The  porous  charcoal  and  gas,  arsenious  sulfide  and 
water,  oil,  soap  and  water,  etc.,  may  all  be  called  disperse 
systems.  They  are  also  called  colloids  and  their  study  colloid 
chemistry.  Even  this  brief  treatment  should  be  sufficient  to 
show  the  reader  the  enormous  importance  of  the  subject.  As  a 
matter  of  fact  the  majority  of  practical  applications  of  chem- 
istry involve  disperse  systems.  Students  will  find  the  five 
lectures  written  on  this  subject  for  the  general  pubic  by  Wolfgang 
Ostwald1  well  worth  reading. 

TOstwald  and  Fisher,  Theoretical  and  Applied  Chemistry.  Published  by 
John  Wiley  &  Sons,  New  York. 


CHAPTER  XXIX 

THE  ATMOSPHERE  AND  RELATED  TOPICS 

765.  The  Composition  of  the  Air. — We  have  already  learned 
that  the  most  important  components  of  air  are  nitrogen,  oxygen, 
water  vapor,  and  carbon  dioxide.     In  addition  to  these  four, 
there  is  but  one  other  component  the  proportion  of  which  exceeds 
o.o i  per  cent.     This  is  the  element  argon  (513),  discovered  in 
the  air  in  1894.     The  percentage  by  volume  of  the  five  named 
components  is  given  in  Table  XXXVII. 

TABLE  XXXVII 

PERCENTAGE  COMPOSITION  OF  THE  AIR 
BY  VOLUME 

Nitrogen 77 .10 

Oxygen 20.70 

Argon 0.80 

Carbon  dioxide o .  03 

Water  vapor  (about)        .      .      .  1.35 

Sum 99.98 

It  is  self-evident  that  the  air  must  also  contain  minute  amounts 
of  several  other  gases  and  vapors,  since  these  are  being  poured 
into  the  air  from  numerous  industrial  sources  and  are  also  being 
formed  by  natural  chemical  changes.  Among  the  minor  com- 
ponents of  the  air  there  may  be  mentioned  hydrogen,  methane, 
sulfur  dioxide,  hydrogen  sulfide,  ammonia,  and  nitrogen  tetroxide. 
These  chemically  active  gases  are  continuously  being  removed 
from  the  air  by  various  means,  so  that  they  never  accumulate 
in  appreciable  proportions.  Other  minor  gaseous  components 
of  the  air  will  be  considered  later  (791-799). 

766.  Why  the  Composition  of  the  Air  Remains  Constant. — 
It  is  plain  that  oxygen  h  removed  from  the  air  on  every  hand 
by  the  burning  of  substances  and  by  the  respiration  of  all  animals. 
If  oxygen  were  not  being  constantly  renewed  the  percentage  of 

491 


492  Introduction  to  General  Chemistry 

this  element  would  in  time  steadily  decrease.  But  as  we  have 
learned  (691),  all  growing  plants  take  in  carbon  dioxide  and 
water,  from  which  they  form,  in  addition  to  other  products, 
cellulose,  or  starch,  and  oxygen,  thus: 

6C02+ 5H20  ->  C6HI0Os-f  602 

The  oxygen  supplied  to  the  air  in  this  way  serves  to  keep  the 
percentage  of  this  element  fairly  constant. 

Closely  connected  with  the  oxygen  balance  in  the  air  is  that 
of  carbon  dioxide.  The  removal  of  this  gas  by  growing  plants 
is  compensated  by  its  formation  in  the  oxidation  of  wood,  coal, 
fuel  gas,  and  other  carbon  compounds  (356-365).  One  other 
agency  has  been  of  much  importance  in  past  geological  ages  in 
diminishing  the  carbon  dioxide  content  of  the  air.  This  is  the 
formation  of  the  shells  of  aquatic  animals,  consisting  largely  of 
calcium  carbonate.  Immense  deposits  of  limestone  (150)  have 
been  built  up  from  the  shells  of  marine  animals. 

The  water- vapor  content  of  the  air  varies  greatly  from  place 
to  place  and  from  time  to  time  in  a  given  place.  Over  the  ocean 
the  lower  layers  of  the  air  are  nearly  saturated  with  water  vapor, 
so  that  the  vapor  pressure  tends  to  approach,  at  each  prevailing 
temperature,  the  value  shown  in  Table  VII  (112).  In  desert 
regions  the  relative  degree  of  saturation  (the  humidity)  is  very 
small. 

The  percentages  of  nitrogen,  oxygen,  and  argon  in  air  freed 
from  water  vapor  and  carbon  dioxide  are  practically  constant  at 
all  times  the  world  over.  The  carbon  dioxide  content  rarely  falls 
below  o .  03  per  cent  but  may  reach  or  exceed  o .  04  per  cent  in 
congested  parts  of  cities.  Indoors  it  may  at  times  run  much 
higher. 

767.  Dew  and  Frost. — After  sundown  the  temperature  of 
the  air  falls,  on  clear  nights  in  particular,  because  of  the  more 
rapid  radiation  of  heat  in  the  absence  of  clouds.  If  the  humidity 
has  been  rather  high  by  day  the  air  may  become  supersaturated 
with  water  vapor  at  night.  For  every  temperature  the  vapor 
of  liquid  water  exerts  a  definite  pressure,  its  saturation  vapor 
pressure  (112,  Table  VII).  If  air  containing  a  fixed  proportion 


The  Atmosphere  and  Related  Topics 


493 


of  vapor  falls  in  temperature  below  the  point  at  which  it  becomes 
saturated,  the  excess  water  separates  out  on  all  exposed  objects 
as  dew.  The  dew  point  is  denned  as  the  temperature  to  which 
a  given  sample  of  moist  air  must  be  cooled  just  to  reach  a  condi- 
tion of  saturation.  If  cooled  further  it  forms  dew. 

When  the  dew  point  lies  below  zero,  if  the  temperature  falls 
sufficiently  low,  the  water  vapor  deposits  in  the  form  of  frost 
(ice). 

768.  Dust  in  the  Air  and  Cloud  Formation. — The  presence 
of  dust  in  the  air  is,  in  general,  considered  nothing  less  than  a 
nuisance;  but,  as  we  shall  point  out  shortly,  it  has  a  very 
important  function  in  connection  with  the  formation  of  clouds 
and  consequently  of  rain.  For  this  reason  especially  we  shall 
consider  the  dust  content  of  the  air. 


FIG.  108 

A  cloud,  or  a  fog  which  is  a  cloud  at  the  earth's  surface,  is 
made  up  of  countless  drops  of  water,  each  so  minute  that  it 
does  not  fall  with  appreciable  speed  (compare  Millikan's  experi- 
ment, 467).  A  cloud  may  form  when  a  mass  of  moist  air  is 
cooled  below  its  dew  point.  But  this  is  true  only  under  certain 
conditions,  as  may  be  demonstrated  by  a  lecture-table  experi- 
ment. If  air  not  specially  freed  from  dust,  and  saturated  with 
water  vapor,  is  suddenly  cooled  several  degrees,  a  cloud  or  fog 
forms  at  once;  but  if  the  air  is  entirely  free  from  dust  no  cloud 
is  formed.  To  show  these  phenomena  experimentally  we  may 
fill  a  large  bottle  (Fig.  108),  with  air  that  has  been  bubbled 


494  Introduction  to  General  Chemistry 

slowly  through  water  in  order  to  insure  its  saturation  with 
vapor.  The  bottle  should  be  closed  with  a  stopper  through  which 
passes  a  glass  tube  attached  by  a  rubber  tube  to  a  second  bottle 
half  the  size  of  the  first.  The  second  bottle  should  have  a  two- 
hole  stopper  and  a  second  tube  leading  to  a  vacuum  pump.  The 
rubber  tube  connecting  the  two  bottles  is  now  closed  by  means 
of  a  clamp  and  the  smaller  bottle  evacuated.  If  now  the  clamp 
is  released,  the  air  in  the  large  bottle  expands  rapidly,  partly 
passing  over  into  the  smaller  bottle,  and  at  the  same  time  a 
dense  cloud  develops  in  the  larger  bottle.  The  expanding  air 
does  work  in  the  process,  and  the  energy  to  do  the  work  is  taken 
from  the  air  as  heat;  in  consequence  the  temperature  falls  so 
much  that  the  air  becomes  supersaturated  at  the  lower  tempera- 
ture, and  the  excess  moisture  separates  as  a  cloud.  If  we  repeat 
this  experiment  with  the  single  change  that  the  air  used  in  filling 
the  large  bottle  is  freed  from  dust  before  it  is  bubbled  through 
water,  by  causing  it  to  filter  through  a  cotton  plug  (A,  Fig.  108), 
no  cloud  forms  upon  expansion.  If  now  a  small  amount  of  dust 
(for  example  a  little  cigarette  smoke)  is  admitted  to  the  partially 
evacuated  large  bottle  a  cloud  is  formed.  (The  smoke  must  be 
introduced  quickly,  before  the  temperature  has  time  to  rise.) 

769.  Counting  Dust  Particles. — Anyone  who  has  noticed  the 
particles  of  dust  floating  in  air,  which  may  be  seen  when  a 
bright  beam  of  light  enters  a  darkened  room,  would  be  inclined 
to  say  that  counting  the  stars  or  the  grains  of  sand  by  the  seaside 
would  be  a  simple  task  in  comparison  with  the  enumeration 
of  the  dust  particles  in  a  given  volume  of  air.  But  thanks  to 
the  facts  related  in  the  preceding  section  the  undertaking  proved 
rather  easy.  Without  dust,  cloud  formation  does  not  take 
place.  In  the  presence  of  dust  each  dust  particle  acts  as  a 
nucleus  for  the  formation  of  a  water  drop.  Therefore  there  are 
as  many  water  drops  in  the  cloud  as  there  were  dust  particles 
present!  A  practical  instrument  has  been  devised  by  Aitken 
to  count,  by  the  aid  of  a  microscope,  the  drops  formed  in  a  known 
volume  of  air  and  thus  to  ascertain  the  dust  content  of  the  air. 
The  number  of  dust  particles  per  c.c.  of  air  varies  greatly  with 
circumstances,  as  would  be  expected.  The  smallest  values,  a 


The  Atmosphere  and  Related  Topics  495 

few  hundred  particles  per  c.c.,  are  found  in  air  above  mid-ocean. 
Clear  mountain  air  contains  a  few  thousand  particles  per  c.c., 
while  the  air  of  large  cities  often  contains  over  100,000  per  c.c. 
In  his  interesting  article  in  the  Encyclopaedia  Britannica  on 
"Dust,"  Aitken  writes: 

Without  atmospheric  dust  not  only  would  we  not  have  the  glorious 
cloud  scenery  we  at  present  enjoy,  but  we  should  have  no  haze  in  the 
atmosphere,  none  of  the  atmospheric  effects  that  delight  the  artist.  The 
white  haze,  the  blue  haze,  the  tender  sunset  glows  of  red,  orange  and  yellow, 
would  all  be  absent,  and  the  moment  the  sun  dipped  below  the  horizon 
the  earth  would  be  in  darkness;  no  twilight,  no  after  glows;  none  of  the 
poetry  of  eventide.  Why  it  may  be  asked  is  this  so  ?  Simply  because  all 
these  are  due  to  matter  suspended  in  the  air,  to  dust. 

770.  The  lonization  of  Gases. — The  topic  about  to  be  dis- 
cussed is  related  to  cloud  formation  in  an  important  way,  and 
for  this  reason  we  shall  have  to  digress  somewhat  before  going 
on  with  our  subject.  It  is  well  known  that  air  under  ordinary 
conditions  is  an  almost  perfect  electrical  insulator.  However, 
it  may  be  made  appreciably  conducting  in 
several  ways.  The  phenomena  may  be 
readily  shown  by  the  use  of  a  gold-  or 
aluminum-leaf  electroscope,  Fig.  109.  If 
the  electroscope  has  good  irisulation  (sulfur 
or  amber,  A,  Fig.  109)  it  will  retain  its 
charge  for  a  long  time.  If  a  flame  of  any 
kind  is  brought  near  the  electrode,  B,  the  FlG 

gold  leaf,  C,  drops  almost  as  rapidly  as  if 
one  had  touched  B  with  the  finger.  The  gases  in  and  about 
a  flame  conduct  electricity  thousands  of  times  better  than 
ordinary  air.  Elaborate  investigations,  which  cannot  profitably 
be  discussed  here,  have  proved  that  part  of  the  gas  molecules 
of  a  flame  are  electrically  charged,  half  positively,  half  negatively. 
Consequently  the  gas  is  said  to  be  ionized;  and  just  as  in  the 
case  of  an  ionized  solution  the  gas  conducts  electricity. 

Every  gas  can  be  ionized  by  a  variety  of  means,  among  which 
are  intense  heat,  X-rays  (476),  and  radium  rays  (480).  Even 
the  most  energetic  means  do  not  convert  into  ions  more  than  a 
very  small  fraction  of  the  total  number  of  gas  molecules. 


496  Introduction  to  General  Chemistry 

771.  Gaseous  Ions  and   Cloud  Formation. — Gaseous  ions 
behave  like  dust  particles  in  being  able  to  serve  as  nuclei  for 
water  drops  in  cloud  formation.     If  dust -free  air  saturated  with 
water  vapor,  is  ionized  by  X-rays,  by  radium,  or  in  any  other 
way  and  is  then  suddenly  cooled  (as  by  expansion  in  the  manner 
described  in  768)  a  cloud  will  form,  each  ion  acting  as  a  nucleus 
for  a  single  drop  of  water. 

772.  The  Formation  of  Rain. — There  are  always  present  in 
the  air  a  small  number  per  c.c.  of  gaseous  ions  formed  in  part  at 
least  by  radioactive  matter  in  the  air.    These  ions,  together 
with  a  much  larger  number  of  dust  particles,  serve  as  the  nuclei 
of  the  drops  forming  ordinary  clouds.     Cloud  formation  occurs 
when  a  current  of  warm,  moist  air  meets  a  current  of  cold  air. 
If  the  water  drops  are  large  enough  they  tend  .to  coalesce  and 
thus  grow  so  large  that  they  fall  to  earth  as  rain.     Rain  is  fre- 
quently accompanied  by  lightning,  the  cause  of  which  we  may 
now  consider. 

773.  The  Cause  of  Lightning. — If  air  is  only  slightly  super- 
saturated with  water  vapor  the  negative  ions  present  are  much 
more  effective  in  the  condensation  of  moisture  than  are  the 
positive  ions,  so  that  it  frequently  happens  that  only  the  nega- 
tive ions  are  at  first  removed  from-  the  upper  layers  of  air  and 
carried  to  the  earth  with  the  falling  rain.    This  results  in  the 
accumulation  of  opposite  electrical  charges  in  the  air  and  on 
the  earth  beneath.     The  lightning  that  frequently  accompanies 
rain  is  the  electric  discharge  between  air  and  earth,  tending  to 
neutralize  the  unlike  charges. 

774.  The  Atmosphere,  a  Disperse  System. — The  behavior 
of  tiny  water  and  dust  particles  in  the  air  suggests  the  behavior 
of  small  particles  in  liquids.    We  have  already  noted  (703)  the 
Brownian  movement  of  smoke  particles.    The  accumulation  of 
an  electrical  charge  by  clouds  in  the  process  of  formation  and 
the  formation  of  rain  with  the  discharge  of  the  former  have  their 
counterparts  in  the  behavior  of  suspensoids.    The  atmosphere 
with  its  dust  and  water  particles  is,  of  course,  a  disperse  system. 

775.  The  Liquefaction  of  Gases. — Faraday's  experiments  on 
the  liquefaction  of  chlorine  have  already  been  described  (242). 


The  Atmosphere  and  Related  Topics 


497 


We  have  also  seen  that  ammonia  can  be  liquefied  when  strongly 
compressed  (517).  Carbon  dioxide  gas  can  be  liquefied  (633) 
at  a  very  high  pressure  if  at  the  same  time  the  temperature  is 
below  31°.  Above  this  temperature  no  pressure,  however  great, 
causes  liquefaction.  At  31°  the  pressure  required  is  72  atmos- 
pheres. At  lower  temperatures  less  pressure  is  required. 
Experiment  has  shown  that  for  every  gas  there  is  a  definite 
temperature  above  which  no  pressure  (however  great)  will  cause 
liquefaction.  This  temperature  is  called  the  critical  tempera- 
ture of  the  gas.  The  vapor  pressure  exerted  at  its  critical 
temperature  by  a  liquefied  gas  is  called  its  critical  pressure. 
Any  gas  may  be  liquefied  if  cooled  below  its  critical  temperature 


TABLE  XXXVIII 


Gas 

Critical  Temperature 
in  Degrees  C. 

Critical  Pressure 
in  Atmospheres 

Boiling-Point  in 
Degrees  C. 

Ammonia  

131° 

H3 

-  39° 

Carbon  monoxide  

-140 

36 

-190 

Carbon  dioxide  

3T 

72 

-   76 

Hydrogen  chloride  
Nitric  oxide  

52 

-  94 

84 
71 

-  35 
-154 

Nitrous  oxide  

37 

72 

-  88 

Sulfur  dioxide  

155 

79 

—    10 

Chlorine  

146 

84 

-  34 

Hydrogen  
Nitrogen  

-243 
-145 

ii 

34 

-252 
-194 

Oxygen 

—  no 

ci 

-182 

and  subjected  to  a  pressure  which  need  not  exceed  its  critical  pressure. 
All  known  gases  except  five  (hydrogen,  nitrogen,  oxygen,  nitric 
oxide,  and  carbon  monoxide)  were  liquefied  by  Faraday  by  the 
year  1845.  The  critical  temperatures,  critical  pressures,  and 
boiling-points  at  one  atmosphere  pressure  of  a  number  of  gases 
are  given  in  Table  XXXVIII. 

Solid  carbon  dioxide  (633)  mixed  with  ether  (642)  to  make 
it  a  better  heat  conductor  was  used  by  Faraday  as  a  cooling 
agent.  By  means  of  this  mixture  a  temperature  of  —80°  is 
easily  obtained.  Even  at  —80°  and  at  very  high  pressures 
Faraday  could  not  liquefy  the  five  named  gases.  These  he 
called  permanent  gases.  We  now  know  that  these  gases  differ 
from  others  only  in  having  critical  temperatures  considerably 


Introduction  to  General  Chemistry 


lower  than  —80°.     They  have  all  been  liquefied  in  more  recent 
times. 

776.  Liquefaction  of  Air. — Practical  methods  for  the  liquefac- 
tion of  air  have  been  devised  by  Hampson,  by  Linde,  and  by 
Claude.  The  principle  of  the  apparatus  of  the  first  two  of  these 
inventors  is  the  same  and  will  be  understood  by  reference  to 
the  diagrammatic  Fig.  no  and  the  following  description.  Air 
is  compressed  to  200  atmospheres  pressure  by  a  powerful  pump, 
A .  The  work  of  compression  is  largely  changed  to  heat  so  that 
the  air  becomes  very  hot.  It  is  next  cooled  by  means  of  running 
water  in  a  cooler,  B,  and  afterwards  passes  through  a  thick- 
walled  steel  cylinder,  C,  containing  solid  caustic  potash  to  free 
it  from  water  vapor  and  carbon  dioxide.  The  compressed  air 


FIG.  no 

next  passes  through  a  metal  pipe  ending  in  a  needle  valve,  D, 
from  which  it  escapes  and  expands  to  normal  atmospheric 
pressure.  Upon  expansion  at  the  needle  valve  the  air  falls 
greatly  in  temperature  (compare  the  preparation  of  carbon 
dioxide  snow,  633),  but  not  sufficiently  to  cool  it  to  its  tempera- 
ture of  liquefaction.  But  the  very  cold  expanded  air  now  flows 
upward  through  the  larger  pipe  surrounding  the  pipe  leading 
to  the  needle  valve  and  thus  cools  the  compressed  air  greatly 
before  it  escapes.  The  result  is  that  the  air  reaching  the  valve 
becomes  colder  and  colder  until  finally  it  reaches  its  liquefaction 
point;  that  is,  its  boiling- temperature  at  atmospheric  pressure. 
From  this  time  on  liquid  air  collects  in  the  lower  part  of  E  and 
can  be  drawn  off  through  a  second  valve. 

777.  The  Properties  of  Liquid  Air. — Liquid  air  is  a  faintly 
blue  mobile  liquid.    It  can  be  kept  in  the  liquid  state  at  atmos- 


The  Atmosphere  and  Related  Topics  499 

pheric  pressure  only  so  long  as  its  temperature  remains  at  or 
below  its  boiling-point,  about  — 190°.  If  it  is  contained  in  an 
open  vessel,  it  boils  away  more  or  less  rapidly,  and  the  absorp- 
tion of  heat  during  its  evaporation  (115)  keeps  its  temperature 
down  to  its  boiling-point  at  atmospheric  pressure.  The  more 
slowly  external  heat  flows  into  the  liquid  air  the  more  slowly  it 
boils  away  and  therefore  the  longer  it  may  be  kept.  Sir  James 
Dewar  has  invented  ingenious  containers  for  liquid  air  that 
provide  the  best  attainable  heat  insulation.  These  are  glass 
vessels  with  double  walls,  having  the  space  between 
the  walls  evacuated  (Fig.  in).  Often  the  walls  are 
silvered.  Glass  is  a  poor  conductor  of  heat,  and  as 
there  is  no  air  between  the  walls  to  transfer  heat  by 
convection,  heat  can  reach  the  interior  only  by  con- 
duction, by  way  of  a  long  path  through  glass. 
Radiant  heat  is  kept  out  by  the  silvered  surface, 
which  reflects  both  heat  and  light.  Heat  reaches  FlG 
the  interior  so  slowly  that  a  five-liter  flask  of  liquid 
air  will  lose  by  evaporation  only  20  to  25  per  cent  of  its  contents 
in  twenty-four  hours.  In  recent  years  Dewar  vessels  have  come 
into  extensive  general  use  (as  "Thermos"  bottles,  "Icy-Hot" 
bottles,  etc.). 

Liquid  air,  being  a  mixture  of  liquid  nitrogen  and  liquid 
oxygen,  does  not  have  a  constant  boiling-point.  Liquid  nitrogen 
boils  at  — 194°  and  liquid  oxygen  at  — 182°.  The  former,  being 
the  more  volatile,  boils  away  more  rapidly,  so  that  the  partially 
evaporated  liquid  is  largely  oxygen.  This  is  decidedly  blue  in 
color.  Liquid  nitrogen  is  colorless.  By  a  process  of  systematic 
fractional  distillation  the  component  gases  of  liquid  air  can  be 
nearly  completely  separated.  These  separate  gases  are  made 
commercially  on  a  large  scale  in  this  way  (309,  513). 

778.  Liquefaction  of  Other  So-called  Permanent  Gases. — 
With  the  liquefaction  of  air  it  became  apparent  that  there  were 
no  permanent  gases  in  the  sense  that  these  could  not  be  liquefied. 
It  was  plain  that  any  gas  could  be  liquefied  either  at  atmospheric 
or  higher  pressure,  provided  it  could  be  cooled  below  its  critical 
temperature.  Faraday  had  failed  to  liquefy  five  gases  because 


500  Introduction  to  General  Chemistry 

he  had  no  means  of  cooling  them  to  sufficiently  low  temperatures. 
With  liquid  air,  boiling  at  atmospheric  pressure,  a  temperature 
of  —190°  is  available.  Since  carbon  monoxide  boils  at  —190° 
and  nitric  oxide  at  —154°  both  these  gases  are  readily  liquefied 
by  cooling  them  with  liquid  air.  In  the  case  of  the  former  a 
little  pressure  above  atmospheric  is  required.  Hydrogen,  how- 
ever, cannot  be  liquefied  at  —190°  even  when  strongly  com- 
pressed. This  fact  indicates  that  "the  critical  temperature  of 
this  gas  is  lower  than  —190°. 

779.  Liquid  Hydrogen. — Hydrogen  was  condensed  to  a 
liquid  (in  appreciable  amounts)  by  Dewar  in  1898.  The  method 
employed  was  similar  to  that  used  to  make  liquid  air  (776); 
but  the  compressed  hydrogen  before  entering  the  liquefier,  D, 
Fig.  no  (776),  was  cooled  to  —185°  by  liquid  air.  Liquid 
hydrogen  is  a  colorless  liquid,  about  one-seventh  as  dense  as 
water.  Its  critical  temperature  is  —243°,  its  critical  pressure 
n  atmospheres.  It  boils  at  —252°  at  atmospheric  pressure, 
equal  to  21°  absolute. 

780.  Flames. — That  a  flame  is  a  burning  gas  is  of  course  well 
known  to  the  reader;  but  it  may  not  have  occurred  to  him  that 
the  flame  of  a  candle  or  kerosene  lamp  is  a  gas  flame.     In  the 
case  of  a  burning  candle  the  wax  melted  by  the  heat  forms  a 
small  pool  of  liquid;  this  the  wick  takes  up  by  capillary  action 
and  brings  to  the  center  of  the  flame,  where  the 
intense  heat  decomposes  the  wax  into  volatile 
(gaseous)  products.    By  holding  a  narrow  glass 
tube  3  inches  long  in  a  candle  flame  so  that  the 
lower  end  of  the  tube  is  at  the  tip  of  the  wick, 
unburned  gas  may  be  drawn  from  the  center  of 
the  flame  and  burned  at  the  upper  end  of  the 
FIG.  1 1 2  tube,  Fig.  112.    This  experiment  also  shows  that 

in  the  center  of  the  candle  flame  the  gas  is  as 
yet  unburned.  It  unites  with  oxygen  in  the  outer  layer  of  the 
flame;  this  is  therefore  the  hottest  part  of  the  flame,  while  the 
interior  is  much  cooler.  If  a  piece  of  writing  paper  is  held  for 
a  few  seconds  in  a  candle  flame  at  the  tip  of  the  wick  and 
perpendicularly  to  the  latter,  it  will  be  scorched  in  a  ring,  the 
center  of  which  is  unburned, 


The  Atmosphere  and  Related  Topics 


The  flame  of  burning  wood  or  coal  is  formed  similarly  to 
the  candle  flame  by  reason  of  the  preliminary  conversion  of  the 
fuel  into  gaseous  products. 

781.  The  Bunsen  Burner. — If  we  close  the  air  vents  at  the 
base  of  a  Bunsen  burner  the  ignited  gas  burns  with  a  flame 
resembling  the  candle  flame  in  structure.  The  flame  is  luminous 
and  is  likely  to  be  smoky.  If  we  open  the  air  vents  sufficient 
oxygen  becomes  mixed  with  the  gas  to  cause  much  more  rapid 
burning,  since  now  it  is  not  necessary  for  gas  and  air  to  mix  by 
the  rather  slow  process  of  diffusion  after  the  gas  has  left  the 
burner  tube.  With  the  ideal  adjustment  of  the  air  vents  some- 
what less  oxygen  is  supplied  than  the  total  necessary,  so  that 
some  oxygen  is  taken  from  the  air  around  the  flame.  If  too 
much  air  mixes  with  the  gas  at  the  vents,  the  mixture  burns 
with  such  great  rapidity  that  the  speed  of  ignition  exceeds  the 
speed  with  which  the  mixed  gas  and  air  travel  upward  in  the 
burner  tube,  with  the  result  that  the  flame 
"strikes  back"  and  "burns  at  the  base." 
In  an  improved  form  of  burner  known  as 
the  Meker  burner  (Fig.  113)  the  design  of 
the  tube  and  vents  is  such  that  a  larger 
proportion  of  air  is  taken  in  than  in  the 
ordinary  Bunsen  burner,  and  air  and  gas 
are  more  intimately  mixed  before  reaching 
the  top  of  the  burner.  To  prevent  strik- 
ing back,  the  wide  upper  end  of  the  burner 
tube  is  provided  with  a  metal  grid,  in  the 
narrow  passages  of  which  the  hot  gases  are 
so  greatly  cooled  that  the  downward  speed  of  the  ignition  wave 
no  longer  exceeds  that  of  the  upward-moving  gas  stream. 

The  Bunsen  flame  with  open  vents  consists  of  two  distinct 
parts:  the  inner  cone  of  a  greenish  color  and  the  outer  cone, 
bluish  in  color  and  less  luminous.  The  hottest  part  of  the  flame 
is  found  in  the  center  of  the  outer  cone,  just  above  the  apex  of 
the  inner  cone. 

782.  Luminous  and  Non-luminous  Flames. — The  cause  of 
luminosity  of  a  gas  flame,  such  as  that  from  a  "fish  tail"  burner, 
has  been  the  subject  of  extensive  investigation.  It  is  probable 


1 


FIG.  113 


502  Introduction  to  General  Chemistry 

that  certain  compounds  in  the  gas,  notably  acetylene,  C2H2 
(663),  are  decomposed  by  heat  giving  free  carbon;  and  that  the 
particles  of  the  latter,  being  intensely  heated,  give  out  light. 
Indeed  we  have  only  to  hold  a  cold  object  in  a  luminous  flame  to 
collect  on  it  a  deposit  of  soot  (carbon).  Furthermore  it  is  well 
known  that  if  acetylene  is  strongly  heated  in  the  absence  of  air 
it  dissociates  into  hydrogen  and  carbon.  The  particles  of  carbon 
in  a  luminous  flame  are  finally  more  or  less  completely  burned. 

The  action  of  air  in  rendering  a  Bunsen  flame  non-luminous 
is  said  by  some  chemists  to  be  due  to  more  perfect  combustion. 
But  this  explanation  is  not  quite  sufficient,  since  a  gas  flame  is 
rendered  non-luminous  by  admixture  with  nitrogen  or  carbon 
dioxide  instead  of  with  air.  Probably  these  inert  gases  so  act 
because  the  temperature  of  the  flame  is  reduced  by  the  dilution 
to  a  point  below  that  at  which  the  acetylene,  etc.,  is  decomposed 
before  it  burns.  Air,  being  four-fifths  nitrogen,  must  also  act 
as  a  diluent  in  the  Bunsen  flame. 

We  always  use  the  non-luminous  Bunsen  flame  when  we 
wish  to  heat  anything  most  efficiently;  and  we  might  be  led  to 
conclude  that  a  given  volume  of  gas  produces  more  heat  as  a 
non-luminous  than  as  a  luminous  flame.  This,  however,  is  not 
true;  the  heat  production  is  exactly  the  same  in  the  two  cases,  that 
is,  if  the  combustion  is  complete  (363).  Nevertheless  a  beaker 
of  water  will  be  heated  more  quickly  and  a  piece  of  glass  will 
be  heated  hotter  in  the  non-luminous  flame.  One  reason  is  this : 
the  luminous  flame  radiates  heat,  as  well  as  light,  in  much  greater 
amount  than  does  the  non-luminous  flame,  as  is  easily  proved 
by  holding  the  hand  at  a  distance  of  a  few  inches  on  one  side  of  a 
Bunsen  flame  and  opening  and  closing  the  air  vent.  Most  of 
the  radiated  heat  is  lost  for  practical  heating  purposes.  A 
second  reason  why  the  non-luminous  flame  is  more  effective  is 
that  it  is  more  concentrated  (compact),  and  that  the  gases  are 
in  more  rapid  motion.  For  both  these  reasons  hot  molecules 
strike  the  object  to  be  heated  more  frequently  than  is  the  case 
in  the  luminous  flame  and  so  raise  its  temperature  more  rapidly. 
783.  Reactions  in  the  Flame.  Bead  Tests. — It  is  perhaps 
a  new  point  of  view  to  consider  the  flame  as  a  reagent,  but  that 


The  Atmosphere  and  Related  Topics  503 

it  is  a  valuable  one  may  be  illustrated  by  some  well-known  bead 
tests.  If  a  metaphosphate  bead  is  made  in  the  usual  way  and 
a  tiny  speck  of  a  copper  salt  is  melted  into  it  by  heating  the  two 
in  the  non-luminous  flame  of  the  burner,  after  the  resulting  bead 
is  cool  it  will  be  found  to  be  blue,  the  usual  color  of  dilute  solu- 
tions of  copper  salts.  But  if  the  air  holes  of  the  burner  are  partly 
closed,  so  that  a  small,  luminous  sheath  appears  on  the  tip  of 
the  inner  cone  of  the  flame  and  the  bead  is  first  melted  in  this 
sheath  and  then  lowered  into  the  inner  cone  of  unburned  gas 
until  it  is  cool,  upon  removal  from  the  flame  the  bead  will  be 
seen  to  be  opaque  and  reddish,  owing  to  the  presence  of  finely 
divided  copper.  If  the  red  bead  is  reheated  in  the  outer  flame 
the  blue  color  reappears.  As  a  consequence  of  such  reactions 
the  inner  cone  of  the  flame  is  known  as  the  reducing  region 
and  the  outer  as  the  oxidizing  region.  The  reducing  region  of 
the  non-luminous  flame  is  relatively  thin,  but  its  depth  is  in- 
creased by  shutting  off  a  little  of  the  air  as  indicated  above. 
Many  other  bead  tests  can  be  made  in  a  similar  way. 

784.  Colored  Flames. — A  number  of  elements  if  brought  into 
a  non-luminous  Bunsen  flame  in  the  form  of  volatile  compounds 

TABLE  XXXIX 

Element  Color  of  Flame 

Sodium Yellow 

Potassium  .                 ...  Violet 

Lithium Crimson 

Calcium Orange 

Strontium Red 

Barium Green 

Copper Bluish  green 

Boron Green 

(usually  salts)  give  to  the  flame  characteristic  colors.  Thus, 
common  salt  and  other  compounds  of  sodium  give  intensely 
yellow  flames.  The  flame  color  is,  in  general,  dependent  on  the 
elements  present  irrespective  of  whether  they  are  free  or  com- 
bined. Table  XXXIX  gives  a  list  of  the  commoner  elements 
forming  colored  flames. 


5°4 


Introduction  to  General  Chemistry 


FIG.  114 


Very  interesting  and  important  facts  about  colored  flames 
are  brought  out  by  the  use  of  the  spectroscope,  to  the  descrip- 
tion of  which  we  shall  now  turn. 

785.  The  Spectroscope. — When  a  narrow  beam,  A,  of  white 
light,  with  parallel  rays,  strikes  a  prism,  B,  in  the  manner  illus- 
trated in   Fig.  114,  its 
colors  are   refracted  to 
different  degrees  so  that 
the    emerging    light    is 
spread  out  as  a   spec- 
trum at  C,  with  its  red 

rays  least  and  its  violet 

rays  most  changed  in 
direction.  A  spectro- 
scope (Fig.  115)  consists  of  a  prism,  A,  with  one  tube,  B, 
carrying  lenses  so  arranged  as  to  throw  a  narrow  beam  of  light 
on  the  prism,  and  a  second  tube,  C,  carrying  a  set  of  lenses 
forming  a  microscope  through  which  the  spectrum  is  viewed.  In 
most  cases  there  is  in  addition  a  third  tube,  Z),  carrying  a  scale 
which  can  be  illuminated,  by  means  of  which  the  different  colors 
of  the  spectrum  can  be  located.  If  one  examines  with  the 
spectroscope  the  light  from  an  incandescent  electric  bulb  he  sees 
the  entire  spectrum  from 
red  on  the  one  hand  to 
violet  on  the  other,  with 
no  single  color  noticeably 
brighter  than  any  other. 
The  same  kind  of  spectrum 
is  given  by  any  white-hot 
body,  as  for  example  an  FIG.  115 

incandescent   gas  mantle. 

Such  a  spectrum  is  called  a  continuous  spectrum.    A  luminous 
gas  flame  also  gives  a  continuous  spectrum. 

786.  Bright-Line  Spectra. — An  entirely  different  picture  is 
presented  when  one  looks  at  a  non-luminous  Bunsen  flame  made 
yellow  by  a  sodium  salt.    Instead  of  the  entire  spectrum  only  a 
bright-yellow  line  appears.    If  the  instrument  is  a  good  one 


The  Atmosphere  and  Related  Topics  505 

and  a  very  narrow  beam  of  light  is  used,  the  yellow  line  is  found 
to  consist  of  two  parallel  lines  very  close  together  in  the  position 
on  the  scale  occupied  by  the  yellow  part  of  the  spectrum  when 
the  entire  spectrum  is  present.  The  spectrum  of  the  element 
lithium  is  still  more  striking,  in  that  it  appears  as  a  single  bright 
line  of  the  purest  red.  The  spectrum  of  potassium  shows  two 
lines,  one  in  the  red  but  not  in  the  same  position  as  that  of  lithium 
or  of  the  same  shade  of  red,  and  another  in  the  violet.  Some 
elements  like  calcium  give  a  rather  complex  spectrum  made  of 
several  lines,  some  of  which  are  rather  broad.  Each  element 
gives  its  own  characteristic  spectrum,  so  that  if  one  is  familiar 
with  the  various  line  spectra  of  the  elements  it  becomes  a  very 
simple  matter  to  identify  at  once  any  element  which  gives  a 
colored  flame.  This  identification  is  made  easy  because  each 
line  of  the  spectrum  always  appears  at  a  fixed  position  on  the 
scale  of  the  instrument,  so  that  the  observer  has  only  to  note  the 
scale  position  of  the  lines  without  considering  critically  their 
colors.  It  is  not  difficult  to  detect  spectroscopically  the  presence 
of  two  or  more  elements  in  a  mixture.  Thus  for  example  if  a 
little  lithium  chloride,  LiCl,  is  mixed  with  some  common  salt 
the  presence  of  the  former  can  easily  be  detected  with  the  spectro- 
scope. To  do  this  the  end  of  a  platinum  wire  is  dipped  in  the 
solution  to  be  tested  and  then  held  in  a  non-luminous  flame 
toward  which  the  slit  of  the  spectroscope  is  directed.  One  sees 
the  lines  of  both  lithium  and  sodium. 

787.  Other  Means  of  Examining  Spectra. — Only  a  small 
number  of  elements  give  colored  flames  suitable  for  spectro- 
scopic  study.  By  suitable  means  every  element  can  be  made  to 
show  a  bright  line  spectrum.  In  general  it  is  necessary  for  this 
purpose  to  heat  the  element  to  a  higher  temperature  than  that 
of  the  Bunsen  flame.  To  do  this  we  may  make  use  of  the  electric 
arc;  or  we  may  cause  sparks  to  pass  between  a  platinum  wire 
and  a  solution  of  the  substance.  In  either  case  bright-line 
spectra  are  seen;  but  there  is  usually  more  or  less  difference 
between  the  flame,  arc,  and  spark  spectra  of  a  given  element.  A 
still  different  method  applicable  to  gases  and  volatile  substances 
consists  in  rendering  them  luminous  by  the  discharge  of  an 


506  Introduction  to  General  Chemistry 

induction  coil  while  they  are  contained  in  a  glass  tube  under 
low  pressure.  The  spectroscope  is  of  great  importance  in  the 
analysis  of  substances,  and  in  several  cases  its  use  has  led  to  the 
discovery  of  new  elements. 

788.  The  Ether  Wave  Hypothesis. — Light  is  a  vibratory 
disturbance  in  the  so-called  luminous  ether,  the  waves  being 
set  up  (according  to  hypothesis)  by  atomic  vibrations.     Light 
of  a  definite  color  is  due  to  waves  all  of  equal  length.     In  so  far 
as  the  visible  spectrum  is  concerned,  red  light  has  the  longest 
waves  and  violet  light  the  shortest.     Waves  of  intermediate 
lengths  give  all  the  intervening  colors  of  the  spectrum.    An 
incandescent  solid  sends  out  waves  of  all  lengths  and  gives 
therefore  a  continuous  spectrum.     On  the  other  hand,  a  glowing 
gas  emits  waves'of  but  very  few  wave-lengths,  corresponding  to 
the  bright  lines  of  its  spectrum.     The  reason  for  this  is  thought 
to  be  that  each  distinct  length  of  wave  is  caused  by  a  single  sort 
of  vibrator.     Very  likely  it  is  the  electrons  composing  an  atom 
(482)  which  act  as  the  vibrators  which  set  up  the  ether  waves 
appearing  to  the  eye  as  light. 

789.  Dark-line  Spectra. — If  a  colored  flame  (say  of  sodium 
light)  is  interposed  between  a  highly  luminous  solid  and  a  spec- 
troscope one  might  expect  to  see  a  continuous  spectrum  with 
its  yellow  portion  crossed  by  a  still  brighter  yellow  line.    The 
effect,  however,  is  quite  the  contrary:  where  we  should  expect 
the  bright  line  of  sodium  a  dark  line  appears  instead.    Under 
similar  conditions  the  colored  flame  of  any  other  element  would 
in  like  fashion  show  a  dark-line  spectrum  of -that  element,  having 
a  dark  line  corresponding  to  every  bright  line  of  its  ordinary 
spectrum.    The  explanation  of  this  curious  fact  seems  to  be  the 
following:    vibrators  (atoms  or  electrons)  which  can  emit  light 
of  a  certain  wave-length  are  also  able  to  absorb  light  of  this  same 
wave-length.    Therefore  those  wave-lengths  of  the  light  of  the 
bright  continuous  spectrum  which  correspond  to  the  bright  lines 
of  the  glowing  gas  are  absorbed  in  passing  through  the  latter, 
so  that  in  their  places  dark  lines  appear  in  the  spectrum. 

790.  The  Composition  of  the  Sun  and  the  Stars. — The  spec- 
trum of  daylight  (sunlight)  shows  a  great  number  of  dark  lines. 


The  Atmosphere  and  Related  Topics  507 

These  were  discovered  by  Wollaston  in  1802  and  first  mapped 
by  Fraunhofer  in  1814;  they  are  commonly  called  Fraunhofer 
lines.  These  lines  correspond  to  the  elements  present  in  the 
sun's  glowing  atmosphere.  Among  the  elements  so  indicated 
are  many  common  on  earth  such  as  calcium,  iron,  hydrogen,  and 
sodium.  A  few  lines  are  present,  however,  which  do  not  cor- 
respond to  any  known  elements.  Some  of  these  are  thought  to 
be  due  to  elements  not  as  yet  discovered  on  earth. 

We  may  even  learn  much  of  the  chemical  composition  of  the 
stars,  which  are  in  fact  very  distant  suns,  by  a  study  of  their 
dark-line  spectra. 

791.  The  Discovery  of  Argon. — In  the  year  1890  there  was 
probably  not  a  single  reputable  chemist  in  the  world  who  would 
have  conceded  the  possibility  of  the  existence  in  the  air,  to  the 
extent  of  nearly  i  per  cent,  of  a  hitherto  unknown  gaseous 
element.  The  discovery  of  argon  in  1894  by  Rayleigh  and 
Ramsay  forms,  therefore,  one  of  the  most  interesting  and  sug- 
gestive chapters  of  modern  chemistry.  The  whole  story  as 
related  by  Lord  Rayleigh  himself  in  the  article  on  argon  in  the 
eleventh  edition  of  the  Encyclopaedia  Britannica  is  well  worth 
reading.  Briefly  stated  the  circumstances  leading  to  the  dis- 
covery are  these:  Lord  Rayleigh,  one  of  the  world's  greatest 
physicists,  and  Professor  Ramsey  (later  Sir  William),  a  brilliant 
chemist,  both  Englishmen,  were  engaged  in  a  research  on  the 
determination  of  gas  densities,  in  which  results  of  the  utmost 
accuracy  were  desired.  Nitrogen  was  one  of  the  gases  studied. 
Now,  as  the  student  well  knows,  nitrogen  can  be  prepared  from 
various  nitrogen  compounds  (513)  as  well  as  also  (so  it  was 
thought)  from  the  air  by  the  removal  of  oxygen,  carbon  dioxide, 
and  water  vapor,  and  the  minute  amounts  of  other  known  gases. 
Rayleigh  and  Ramsey  found  that  nitrogen  extracted  in  this  way 
from  air  had  a  density  half  of  i  per  cent  greater  than  nitrogen 
made  from  ammonia  or  other  pure  compounds  of  nitrogen.  This 
difference  in  density  was  about  fifty  times  as  great  as  could  be 
accounted  for  by  experimental  error!  The  only  logical  conclu- 
sion was  that "  atmospheric  nitrogen"  contained  an  appreciable 
proportion  of  a  new  gas  denser  than  nitrogen  and  equally  or 


508  Introduction  to  General  Chemistry 

more  inert  chemically  than  the  latter.  A  thoroughgoing  search 
of  the  literature  of  nitrogen  brought  to  light  a  most  valuable 
clue.  Nearly  a  century  earlier  Cavendish  (292),  the  discoverer 
of  hydrogen,  had  carried  out  an  experiment  in  which  atmospheric 
nitrogen  mixed  with  an  excess  of  oxygen  was  subjected  to  the 
prolonged  action  of  electric  sparks  (566) .  The  oxides  of  nitrogen 
so  formed  were  absorbed  in  alkali,  and  the  volume  of  the  nitrogen 
was  thus  reduced  to  1/120  of  that  taken.  But  further  diminution 
of  volume  could  not  be  made  to  take  place. 

Rayleigh  and  Ramsay's  repetition  of  Cavendish's  experiment 
with  much  more  refined  and  larger  apparatus  enabled  these 
scientists  to  obtain  considerable  amounts  of  the  chemically 
inert  residual  gas.  When  this  was  freed  from  nitrogen  and  other 
known  gases  it  was  found  to  be  a  colorless,  odorless  gas  i  .25 
times  as  dense  as  oxygen  (i  liter  weighs  i .  78  g.).  It  was  named 
argon  and  given  the  symbol  A. 

Every  effort  to  cause  argon  to  combine  with  other  elements 
or  react  in  any  way  with  any  other  substance  proved  completely 
futile.  Furthermore  argon  could  not  be  decomposed  by  any 
physical  means.  It  was  therefore  classed  as  an  element,  but 
one  without  chemical  affinity  and  therefore  without  chemical 
properties. 

792.  The  Molecular  and  Atomic  Weights  of  Argon. — At 
standard  conditions,  22  .4  liters  of  argon  weigh  39.9  g.;  there- 
fore the  molecular  weight  of  the  element  is  39.9.  Since  argon 
does  not  combine  with  other  elements  it  is  impossible,  by 
methods  so  far  discussed  in  this  text,  to  determine  its  atomic 
weight.  If,  however,  we  can  by  some  independent  method  find 
the  number  of  atoms  in  a  molecule  of  argon  we  have  only  to 
divide  its  molecular  weight  by  this  number  to  get  its  atomic 
weight. 

By  means  of  calculations  based  on  the  kinetic-molecular 
hypothesis  it  has  been  found  that  if  no  expansion  is  permitted 
the  amount  of  heat  required  to  raise  the  temperature  of  one  gram 
molecular  weight  of  a  gas  one  degree  (the  so-called  molecular 
heat)  should  be  three  calories,  if  each  molecule  consists  of  a 
single  atom.  In  this  case  the  heat  applied  is  wholly  used 


The  Atmosphere  and  Related  Topics  509 

to  increase  the  velocities  of  the  molecules.  If  however  the 
molecules  are  made  up  of  two  or  more  atoms  each,  then  part 
of  the  heat  is  used  up  in  increasing  the  internal  energy  of  each 
molecule  (that  is,  increasing  the  vibrational  or  rotational 
velocities  of  the  atoms  within  each  molecule  with  respect  to 
one  another).  Therefore  the  heat  required  to  raise  the  tempera- 
ture of  one  gram  molecular  weight  of  a  diatomic  gas  should 
theoretically  be  more  than  three  calories. 

Now  it  has  been  found  by  experiment,  in  strict  accord  with 
theory,  that  the  vapors  of  sodium  and  mercury,  elements  having 
single  atom  molecules  (75),  have  molecular  heats  of  three 
calories.  No  known  gas  or  vapor  has  a  smaller  molecular  heat 
than  this  value.  On  the  other  hand,  just  as  theory  predicts, 
oxygen,  hydrogen,  and  nitrogen,  all  forming  diatomic  mole- 
cules, have  molecular  heats  of  about  4.9  calories.  Gaseous 
elements  and  compounds  with  still  more  complex  molecules  show 
still  larger  molecular  heats.  When  the  molecular  heat  of  argon 
was  found  to  be  three  calories  the  conclusion  was  at  once  drawn 
that  its  molecules  consist  of  single  atoms  and  therefore  that  its 
atomic  weight  is  identical  with  its  molecular  weight,  namely 
39 .9.  Indeed  this  conclusion  is  quite  in  harmony  with  the  fact 
that  argon  does  not  unite  with  other  elements,  for  it  would  be 
very  strange  if  its  atoms  united  with  one  another  when  they 
show  no  inclination  to  unite  with  other  atoms. 

793.  The  Ratio  of  Molecular  Heats  of  Gases. — The  molec- 
ular heat  referred  to  in  the  previous  section  is  the  so-called 
molecular  heat  at  constant  volume,  Cv.  It  represents  the  heat 
necessary  to  raise  one  gram  molecular  weight  of  the  gas  through 
one  degree  centigrade  when  the  volume  of  the  gas  is  kept  con- 
stant. If  the  gas  is  allowed  to  expand  at  constant  pressure 
while  it  is  being  heated,  two  more  calories  per  molar  weight  are 
needed  to  raise  the  temperature  one  degree.  These  two  calories 
are  used  up  in  the  work  of  overcoming  the  external  pressure 
during  the  expansion  of  the  gas.  The  value  so  determined 
is  the  molecular  heat  at  constant  pressure,  CP.  Now  it 
appears  that  the  ratio  of  the  two  molecular  heats  CP/CV  for  a 
given  gas  may  be  calculated  from  the  velocity  of  sound  in  the 


510  Introduction  to  General  Chemistry 

gas  in  question,  and  that  the  ratio  is  a  relatively  easy  value  to 
determine.  As  a  result  it  is  usual  to  find  CP/CV  instead  of  the 
separate  values.  For  if  Cv  for  monatomic  gases  is  equal  to 
three,  then  the  ratio  of  the  molecular  heats  may  be  represented 
thus: 


The  ratio  of  the  molecular  heats  of  a  diatomic  gas  is  in  no  case 
more  than  i  .40.  The  ratio  of  the  molecular  heats  is  at  its  maxi- 
mum in  the  case  of  monatomic  gases. 

794.  The  Story  of  the  Discovery  of  Helium.  —  Of  all  the 
elements  known  in  1894  argon  stood  alone  in  refusing  to  unite 
chemically  with  any  other  element.  Naturally  this  proved  a 
matter  of  surpassing  interest  to  chemists  at  the  time.  After 
the  extensive  laboratory  experiments  in  which  Rayleigh  and 
Ramsay  hoped  to  make  compounds  of  argon  had  all  failed,  it 
occurred  to  these  scientists  that  such  compounds  might  possibly 
exist  in  nature.  But  if  this  were  true,  where  should  they  seek 
these  strange  substances?  Once  more  a  search  of  old  chemical 
literature  furnished  a  clue.  In  a  paper  by  the  American  chemist 
Hillebrand  there  appeared  a  statement  that  the  mineral  cleveite 
when  heated  gave  off  nitrogen.  Before  the  discovery  of  argon 
it  was  considered  sufficient  identification  of  a  gas  as  nitrogen 
to  show  that  it  was  chemically  inert  and  not  one  of  the  other 
known  gases.  Perhaps  the  "nitrogen"  from  cleveite  might  be 
argon!  Ramsay  repeated  Hillebrand  's  experiment  and  obtained 
very  readily  an  inert  gas,  but  it  was  not  nitrogen,  since  it  would 
not  unite  with  oxygen  under  the  action  of  electric  sparks  (566). 
Moreover,  it  was  not  argon,  since  its  density  was  but  one-tenth 
as  great  as  that  of  argon.  It  was  plainly  a  new  element,  with 
a  density  only  twice  that  of  hydrogen;  but  like  argon,  entirely 
devoid  of  chemical  affinity.  The  spectrum  of  the  new  gas  showed 
a  yellow  line  not  far  from  that  of  sodium  but  not  identical  with 
that  of  any  hitherto  known  element.  Search  of  the  literature 
of  spectroscopy,  however,  revealed  a  most  interesting  fact.  In 
1868  the  astronomer  Lockyer  studied  the  photographs  of  spectra 
of  the  sun's  corona  (the  incandescent  atmosphere  of  the  sun) 


The  Atmosphere  and  Related  Topics  511 

taken  during  a  total  eclipse  and  had  noted  the  presence  of  a 
Fraunhofer  dark  line  in  the  yellow  of  the  spectrum  which  could 
not  be  identified  as  due  to  any  known  element.  In  explanation 
Lockyer  suggested  that  this  line  was  doubtless  due  to  a  new 
element  to  which  he  gave  the  name  helium  (Gr.  helios,  the  sun). 
Careful  comparison  of  the  position  of  the  dark  line  of  the  Lockyer 
photograph  with  the  yellow  line  of  the  new  gas  showed  beyond 
any  doubt  that  the  two  were  produced  by  the  same  element. 
Therefore  the  new  gas  was  Lockyer's  helium!  Thus  Rayleigh 
and  Ramsay's  search  for  a  possible  compound  of  argon  led  to  the 
discovery  on  earth  of  an  element  whose  existence  in  the  sun  had 
been  announced  thirty  years  earlier  on  the  evidence  of  a  photo- 
graph taken  of  an  object  92  million  miles  away. 

795.  Helium   and   the   Alpha   Rays   of   Radioactive    Sub- 
stances.— We  have  already  stated  (480)  that  the  alpha  rays  of 
radioactive  substances  are  helium  atoms.     The  further  discus- 
sion of  this  scientifically  important  topic  will  be  taken  up  in 
chapter  xxxi.    At  this  point  we  shall  only  add  that  minerals 
like  cleveite,  in  which  helium  is  contained,  all  contain  radio- 
active substances,  and  that  the  helium  present  is  beyond  doubt 
the  product  of  radioactive  changes  (867).     Helium  is  present  in 
the  atmosphere  in  minute  amount,  about  four  parts  in  a  million. 

796.  The  Properties  of  Helium. — As  we  have  stated,  helium, 
like  argon,  is  devoid  of  all  chemical  properties.     It  is  a  colorless, 
odorless  gas  of  twice  the  density  of  hydrogen.     One  liter  weighs 
o.iSg.  and  22.4  liters  weigh  4.og.     Therefore  its  molecular 
weight  is  4.     Since  the  ratio  of  its  two  molecular  heats  (793)  is 
i .  66,  it  is  a  monatomic  gas  and  its  atomic  weight  is  therefore  4. 
Of  all  gases  helium  is  the  most  difficult  to  liquefy.     It  has  the 
lowest  critical  temperature,  —268°  or  5°  absolute.     By  cooling 
compressed  helium  by  boiling  liquid  hydrogen  and  allowing  it 
to  expand  in  a  liquefier  similar  in  principle  to  that  used  for  liquid 
air  (776)  it  has  been  condensed  to  a  liquid  which  boils  at  4.3° 
absolute.    The  boiling-point  of  helium  under  reduced  pressure 
showed  a  temperature  of  about  3°  absolute;   this  is  up  to  the 
present  time  the  lowest  temperature  ever  produced  experi- 
mentally.   The  liquefaction  of  helium  was  accomplished  by 


512  Introduction  to  General  Chemistry 

Dr.  Kammerlingh  Onnes,  of  the  University  of  Leiden,  in  the 
year  1908. 

797.  Helium  Balloons. — The  use  of  balloons  in  warfare  is 
extremely  precarious  on  account  of  the  inflammability  of  the 
hydrogen  with  which  they  are  filled.  Helium  being  incom- 
bustible would  be  an  ideal  substitute  for  hydrogen  if  it  could  only 
be  had  in  sufficient  quantity.  The  lifting  power  of  a  helium 
balloon  is  eleven-twelfths  of  that  of  the  same  balloon  filled  with 
hydrogen.  For  a  number  of  years  it  had  been  known  that  helium 
was  often  found  in  small  proportions  in  natural  gas,  and  when 
the  matter  of  helium  balloons  began  seriously  to  be  considered 
in  the  last  year  of  the  war  (1918)  it  was  found  that  a  certain 
gas  well  in  Texas  produced  gas  containing  nearly  i  per  cent  of 
helium.  Calculation  showed  that  this  well  also  produced  a 
sufficient  total  amount  of  helium  to  be  of  practical  importance. 
The  United  States  government  appropriated  $500,000  for  a 
plant  to  produce  pure  helium,  and  at  the  end  of  the  war  pro- 
duction had  already  begun.  It  is  not  unlikely  that  immense 
dirigibles  filled  with  helium  will  in  the  near  future  furnish  a 
safe  means  of  aerial  travel.  If  this  turns  out  to  be  the  case  it 
will  be  because  two  English  scientists  were  keen  to  discover  the 
cause  of  a  seemingly  trivial  matter — the  small  discrepancy  in 
the  density  of  nitrogen  as  prepared  in  different  ways!  But  this 
is  the  usual  course  of  great  practical  discoveries,  the  scrupulous 
following  up  of  every  outstanding  scientific  inconsistency. 

798.  Three  Other  Inactive  Gases. — Shortly  after  the  dis- 
covery of  helium  three  other  inert  gases  were  discovered  in 

TABLE  XL 


Gas 

Symbol 

Atomic  Weight 

Neon  

Ne 

2O 

Krypton  

Kr 

83 

Xenon 

X 

1  3O 

minute  amount  in  the  atmosphere.  Like  helium  and  argon 
these  new  gases  do  not  form  any  chemical  compounds;  they  are 
also  monatomic.  These  gases  with  their  symbols  and  atomic 
weights  are  given  in  Table  XL. 


Ttie  Atmosphere  and  Related  Topics  513 

799.  Summary  and  Conclusions. — The  air  may  be  considered 
as  a  great  gaseous  ocean  into  which  flow  many  gases  of  terrestrial 
origin  and  from  which  are  drawn  the  gases  entering  into  the 
earth's  chemical  activity.  It  is  evident  that  in  addition  to  the 
several  well-known  components  of  the  air  there  must  also  be 
present  numerous  other  gases  and  vapors  in  minute  proportions. 
The  composition  of  the  air  remains  nearly  constant  because  of 
the  approximate  equality  of  formation  and  removal  of  the  several 
components.  In  addition  to  the  gases  present,  water  vapor  and 
dust  are  components  of  prime  importance.  Without  dust  (or  gas- 
eous ions)  we  should  have  no  clouds,  and  without  clouds  no  rain, 
although  we  should  have  high  humidity  and  tremendous  dews. 
Viewed  in  the  light  of  the  matters  considered  in  chapter  xxviii 
we  must  see  that  the  atmosphere  is  a  disperse  system  consisting 
of  a  complex  gaseous  mixture  filled  with  a  suspension  of  minute 
solid  particles  (dust)  and  liquid  droplets  (water).  The  gaseous 
ions  of  the  air  become  in  part  attached  to  these  suspended 
particles  giving  to  them  plus  and  minus  charges,  corresponding 
to  the  charges  on  particles  of  colloidal  solutions. 

The  work  on  the  liquefaction  of  gases  so  ably  started  by 
Faraday  in  1823  was  brilliantly  completed  when  Kammerlingh 
Onnes  liquefied  helium  in  1908.  Every  known  gas  has  now  been 
liquefied,  and  all  save  helium  have  been  solidified.  In  the 
liquefaction  of  gases  two  principles  are  most  noteworthy:  first, 
the  gas  can  exist  as  a  liquid  only  below  its  critical  temperature ; 
second,  a  highly  compressed  and  sufficiently  cold  gas  will,  upon 
being  allowed  to  expand,  fall  greatly  in  temperature. 

The  spectra  of  elements  volatilized  at  high  temperatures,  as 
in  a  flame  or  by  means  of  an  electric  spark  or  arc,  appear  as 
bright  lines.  Each  element  gives  definite  and  characteristic 
lines.  If  white  light  from  a  glowing  solid  passes  through  an 
incandescent  gas,  the  continuous  spectrum  of  the  former  is  seen 
to  be  crossed  by  dark  lines  which  are  the  Fraunhofer  spectrum 
of  the  gas.  The  spectroscope  is  of  great  importance  for  the 
detection  of  elements.  By  its  aid  the  composition  of  the  atmos- 
phere of  the  sun  and  of  the  stars  has  been  definitely  revealed. 


CHAPTER  XXX 
SOME  ADDITIONAL  ELEMENTS  AND  THEIR  COMPOUNDS 

800.  Introduction. — The  total  number  of  elements  known  at 
the  present  time  is  83  (if  we  exclude  the  products  of  radioactive 
change,  chap,  xxxii).     These  with  their  symbols  and  atomic 
weights  are  given  in  Table  XLI.     In  the  foregoing  chapters  we 
have  become  acquainted  with  but  little  more  than  one-third  of 
the  elements.     In  the  present  chapter  we  shall  study  more  or 
less  briefly  the  chemistry  of  a  number  of  additional  elements. 
These,  with  those  studied  earlier,  include  all  of  the  elements  of 
practical  importance.     This  list  does  not  embrace  several  ele- 
ments which  are  fairly  abundant  but  for  which  no  technical  uses 
have  as  yet  been  discovered.     In  chapter  xxxi  these  remaining 
elements  will  also  be  considered. 

Our  study  of  part  of  the  elements  has  enabled  us  to  develop 
and  illustrate  many  principles  and  laws;  but  the  student  will 
naturally  wonder  whether  the  study  of  the  remaining  elements 
will  bring  forth  an  entirely  new  set  of  generalizations.  We  can 
hasten  to  assure  him  that  this  is  not  the  case,  for  in  the  study 
of  new  elements  and  compounds  he  will  encounter  very  few 
facts  and  phenomena  that  do  not  have  their  counterparts  among 
those  of  more  familiar  substances. 

80 1.  Boron. — Borax  and  boric  acid  (also  known  as  boracic 
acid)  are  doubtless  known  to  everyone.    These  are  compounds 
of  the  non-metallic  element  boron.    This  element  occurs  plenti- 
fully as  boric  acid,  H3B03,  and  also  as  borates  of  sodium,  magne- 
sium, and  calcium,  but  never  as  free  boron.     It  is  found  chiefly  in 
Italy,  California,  and  Thibet.     Boric  acid,  which  is  volatile  with 
steam,  is  present  in  the  steam  escaping  from  the  earth  in  certain 
volcanic  regions,  especially  in  Italy.    Sodium  borate,  or  borax, 
which  is  easily  soluble,  occurs  abundantly  in  the  waters  of  some 
California  and  Nevada  lakes.     It  is  also  found  in  the  deposits 
left  by  the  drying  up  of  certain  lakes  of  this  same  region.     Borax 

5J4 


Some  Additional  Elements  and  Their  Compounds        515 

is  also  made  technically  from  the  borates  of  magnesium  and 
calcium.     Boric  acid  and  borax  are  the  only  compounds  of  this 


TABLE  XLI 

INTERNATIONAL  ATOMIC  WEIGHTS,  1917 


Element 

Symbol 

Atomic 
Weight 

Element 

Symbol 

Atomic 
Weight 

Aluminum  
Antimony  

Al 
Sb 

27.1 
1  2O    2 

Molybdenum  .... 
Neodymium 

Mo 

Nd 

96.0 
144    3 

Argon  

A 

39  88 

Neon  .  .  . 

Ne 

2O    2 

Arsenic  

As 

74.96 

Nickel  

Ni 

58  68 

Barium 

Ba 

1  37    37 

Niton  (radium 

Beryllium 

Be 

0    I 

emanation) 

Nt 

222    4 

Bismuth  

Bi 

208.0 

Nitrogen  

N 

14  01 

Boron  . 

B 

no 

Osmium 

Os 

I  no  Q 

Bromine. 

Br 

70    Q2 

Oxygen. 

o 

1  6  oo 

Cadmium  

Cd 

112    40 

Palladium 

Pd 

1  06    7 

Caesium  

Cs 

132  81 

Phosphorus 

P 

•5  I      O4 

Calcium  

Ca 

40   07 

Platinum.  . 

Pt 

IQIT    2 

Carbon 

I  2    OO5 

Potassium 

K 

3O    IO 

Cerium  

Ce 

140.  25 

Praseodymium  .  .  . 

Pr 

140  o 

Chlorine 

Cl 

•}f    4.6 

Radium 

Ra 

226  o 

Chromium. 

Cr 

<2    O 

Rhodium 

Rh 

102  9 

Cobalt 

Co 

eg   07 

Rubidium 

Rb 

8c  4.< 

Columbium.  .    .  . 

Cb 

03    I 

Ruthenium 

Ru 

IOI    7 

Copper  
Dysprosium  
Erbium  
Europium  

Cu 
Dy 
Er 
Eu 

63-57 

162.5 
167.7 

1  552  o 

Samarium  
Scandium  
Selenium  
Silicon  

Sa 
Sc 
Se 
Si 

I50-4 
44.1 

79-2 
28  3 

Fluorine  

F 

19.0 

Silver  

Ag 

107.88 

Gadolinium  

Gd 

icy.  7 

Sodium  

Na 

23  .OO 

Gallium 

Ga 

60  o 

Strontium 

Sr 

87    63 

Germanium. 

Ge 

72    S 

Sulfur 

s 

^2    06 

Gold 

Au 

IQ7    2 

Tantalum 

Ta 

181   5 

Helium  

He 

4  oo 

Tellurium  .  .  . 

Te 

127.  ? 

Holmium  

Ho 

163  «; 

Terbium  

Tb 

I^Q.  2 

Hydrogen  

H 

1.008 

Thallium  

Tl 

204.0 

Indium  

In 

114.8 

Thorium  

Th 

232  .4 

Iodine  
Iridium    .  . 

I 
Ir 

126.92 

IQ-2      I 

Thulium  
Tin. 

Tm 
Sn 

168.5 

118  7 

Iron  

Fe 

55     84 

Titanium. 

Ti 

48  I 

Krypton  
Lanthanum  

Kr 
La 

82.92 
I39.O 

Tungsten  
Uranium  

•W 

u 

184.0 
238.2 

Lead 

Pb 

2O7      2O 

Vanadium 

v 

rri    o 

Lithium  

Li 

6-94 

Xenon  

Xe 

130.2 

Lutecium. 

Lu 

175    O 

Ytterbium 

Yb 

173  5 

Magnesium  
Manganese  
Mercury.  . 

Mg 
Mn 
Hg 

24.32 

54-93 
200.  6 

Yttrium  
Zinc  
Zirconium  .  . 

Yt 
Zn 
Zr 

88.7 
65-37 

00.6 

element   that   are   of   any  commercial  importance;    these  are 
cheap  substances  and  are  produced  in  large  quantities. 


516  Introduction  to  General  Chemistry 

802.  Boric  Acid,  H3BO3. — Boric  acid  forms  white  crystals 
which  dissolve  in  25  parts  of  cold  or  3  parts  of  boiling  water. 
It  is  a  very  weak  acid,  weaker  even  than  carbonic  acid.    Its  solu- 
tion is  almost  tasteless;    it  certainly  has  no  sour  taste.    The 
solution  is  used  extensively  in  medicine  as  an  antiseptic  lotion. 
A  cold,  saturated  solution  is  an  excellent  eye  wash.    If  alkali 
has  by  accident  got  into  the  eye,  after  the  former  has  been  washed 
out  at  once  with  water,  boric  acid  solution  should  be  copiously 
applied.    In  the  past  boric  acid  was  rather  extensively  used  as  a 
preservative  for  foods,  especially  milk.    Its  use  in  this  way  is 
now  prohibited,  although  it  is  not  poisonous  if  taken  internally 
in  small  amount.     It  is  sometimes  prescribed  in  medicine  for 
internal  use  (dose  0.3  to  i  g.). 

Boric  acid  is  readily  decomposed  by  heat  into  water  and 
metaboric  acid,  HBO2,  tetraboric  acid,  H2B4O7,  and  finally 
boron  trioxide,  B203: 

H3BO3->HBO2+H2O, 
4HBO2-»H2B407+H2O, 
H2B4O7->2B2O3+H2O . 

The  first  two  reactions  take  place  at  rather  low  temperatures, 
the  last  at  a  red  heat.  These  reactions  are  analogous  to  those 
by  which  pyrophosphoric  and  metaphosphoric  acids  are  formed 
from  orthophosphoric  acid  (590) . 

The  ordinary  salts  of  boric  acid  are  those  derived  from  meta- 
boric, or  more  commonly  tetraboric,  acid.  Borax,  sodium  tetra- 
borate,  Na2B407«ioH2O,  forms  colorless,  glassy  crystals.  Boric 
acid  is  easily  made  from  a  solution  of  borax  by  adding  hydro- 
chloric or  sulfuric  acid  to  a  hot  solution.  On  cooling,  crystals  of 
boric  acid  separate  out: 

Na2B407+2HCl+5H2O^2NaCl+4H3BO3 . 

803.  Borax.— Borax,  or  sodium  tetraborate,  Na2B4O7-ioH2O, 
is  the  only  salt  of  boric  acid  of  commercial  importance.     It  is 
easily  soluble  in  water,  and  its  solution  is  alkaline  by  reason  of 
hydrolysis  (436).     Borax  is  used  extensively  in  the  laundry  for 
softening  hard  water  (156).    Hardness  of  water  is  due  to  the 


Some  Additional  Elements  and  Their  Compounds        5  ry 

presence  of  the  bicarbonates  and  sulphates  of  calcium  princi- 
pally and  magnesium  to  a  smaller  extent.  The  bicarbonates 
cause  temporary  hardness,  so  called  because  they  are  decomposed 
and  the  carbonates  precipitated  when  the  water  is  boiled. 
The  two  sulphates  cause  what  is  called  permanent  hardness. 
Borax  softens  water  by  precipitating  all  the  calcium  and  most  of 
the  magnesium  present  in  hard  water,  as  illustrated  for  the 
calcium  salts  by  the  following  equations : 

Na2B4O7H-Ca(HCO3)2+5H2O^Na2C03+CaCO3+4H3BO3 

CaSO4+Na2CO3->CaCO3+Na2SO4.  (448) 

Upon  being  heated  borax  partially  melts  and  at  the  same 
time  gives  off  its  water  of  hydration  and  in  so  doing  swells  to  a 
spongy,  viscous  mass  (intumesces) ;  this  finally  melts  com- 
pletely when  all  the  water  has  been  driven  off,  and  it  then 
resembles  melted  glass.  The  cold  product,  which  is  anhydrous 
borax,  Na2B407,  is  a  clear,  colorless,  brittle  solid  called  borax 
glass.  Beads  made  of  borax  glass  are  used  in  the  same  way  as 
metaphosphate  beads  (596),  which  they  closely  resemble  in 

TABLE  XLII 
COLORS  OF  BORAX  BEADS 


Element 

Oxidizing  Flame 

Reducing  Flame 

Copper 

Blue 

Red  particles  of  metal 

Chromium 

Green 

Green 

Iron  

Yellow 

Light  green 

Nickel  .  .  . 

Brown 

Gray  particles  of  metal 

Cobalt  
Manganese  

Blue 
Amethyst 

Blue 
Colorless 

appearance.  Various  metallic  oxides  readily  dissolve  in  fused 
borax  glass,  and  in  many  cases  the  beads  have  colors  char- 
acteristic of  the  added  metallic  element.  If  the  element  is 
readily  oxidized  or  reduced  the  color  of  the  bead  will  depend  on 
whether  it  has  been  heated  in  an  oxidizing  or  a  reducing  flame 
(783).  Table  XLII  shows  the  colors  of  borax  beads  of  several 
elements. 


518  Introduction  to  General  Chemistry 

In  explanation  of  the  action  of  melted  borax  on  oxides  it 
may  be  supposed  that  reactions  of  the  type  illustrated  by  the 
following  equation  occur: 

Na2B4O7+CuO-»2NaB02+Cu(B02)2 . 

Elements  which  give  differently  colored  beads  in  the  oxidizing 
and  reducing  flames  do  so  because  they  undergo  oxidation  and 
reduction.  Thus  the  yellow  iron  bead  obtained  in  the  oxidizing 
flame  contains  ferric  metaborate,  while  the  green  one  from  the 
reducing  flame  contains  a  ferrous  salt.  The  borates  of  copper 
and  nickel  are  reduced  to  the  corresponding  metals  in  the 
reducing  flame.  It  may  be  added  that  boron  trioxide  also  forms 
a  glass  after  being  fused,  and  that  it  also  readily  unites  with 
metallic  oxides  at  high  temperatures  to  form  borates.  Some 
useful  optical  glasses  contain  borates  as  essential  ingredients. 
Borates  are  also  used  in  making  some  kinds  of  enamels  and 
glazes.  In  the  welding  of  iron  a  little  borax  sprinkled  on  the 
hot  metal  dissolves  the  iron  oxide  always  present  and  thus  gives 
a  clean  surface  for  the  weld. 

804.  Silicon. — The  element  silicon  is  the  second  most  abun- 
dant of  the  components  of  the  earth's  crust.  It  is  estimated 
that  about  50  per  cent  of  the  known  part  of  the  earth  is  oxygen 
and  25  per  cent  is  silicon.  This  is  not  surprising  when  we  realize 
that  common  quartz  sand  is  essentially  silicon  dioxide,  SiO2,  and 
that  numerous  minerals  composing  the  bulk  of  the  earth  are 
compounds  of  silicon.  Among  such  minerals,  besides  quartz, 
we  may  mention  clay,  granite,  agate,  and  opal  as  being  well 
known.  The  element  silicon  does  not  occur  free  in  nature, 
although  it  can  be  made  by  reduction  of  the  oxide  with  carbon 
at  the  high  temperature  of  the  electric-arc  furnace: 

Si02+2C->Si+2CO . 

Free  silicon  made  in  this  way  is  only  90  to  98  per  cent  pure.  It 
is  a  semimetallic  solid  somewhat  resembling  cast-iron  in  appear- 
ance and  having  a  density  of  2 .35  and  a  melting-point  of  about 
1400°.  It  conducts  electricity  fairly  well.  It  is  not  attacked 
by  any  of  the  common  acids,  with  the  exception  of  hydrofluoric, 


Some  Additional  Elements  and  Their  Compounds         519 

with  which  it  reacts  to  form  silicon_tetrafluoriifi,  SiF4  (270). 
Metallic  silicon,  as  it  is  called  by  technical  men,  is  too  brittle 
to  be  of  much  importance  for  the  manufacture  of  vessels  or 
machines;  but  an  alloy  with  iron  containing  about  14  per  cent 
of  silicon  finds  important  use  in  chemical  industry  for  the  con- 
struction of  vessels,  pumps,  etc.,  for  the  handling  of  strongly 
acid  solutions.  These  alloys,  known  under  the  trade  names  of 
Duriron,  Tantiron,  etc.,  are  extremely  resistant  to  the  action 
of  acids.  Ferro-silicon  is  an  alloy  of  iron  and  silicon  containing 
50  to  75  per  cent  of  the  latter  and  is  used  in  making  other  silicon 
iron  alloys  containing  less  silicon. 

805.  Compounds  of  Silicon  with  Hydrogen  Carbon  and 
Chlorine. — Silicon  forms  a  great  many  compounds  with  hydro- 
gen, carbon,  and  chlorine,  of  much  theoretical  but  little  practical 
importance.  Many  of  these  compounds  are  analogous  to  certain 
compounds  of  carbon.  Thus  we  have  silico-methane,  SiH4,  a 
combustible  gas,  and  silico-ethane,  Si2H6,  a  spontaneously  com- 
bustible liquid,  corresponding  to  methane  and  ethane  (643) 
respectively.  Silicon  tetrachloride,  SiCl4,  is  a  colorless  liquid 
formed  by  the  union  of  silicon  and  chlorine  and  also  by  the  action 
of  chlorine  on  a  red-hot  mixture  of  silica  and  carbon  (charcoal) : 
SiO2+2C+2Cl2->SiCl4+2CO . 

Silicon  tetrachloride  boils  at  58°  and  resembles  in  many  respects 
carbon  tetrachloride,  CC14  (644),  of  boiling-point  76°.  Unlike 
the  latter,  the  former  reacts  readily  with  water.  Silicic  acid, 
H2Si03,  and  HC1  are  formed: 

SiCl4+3H2O->H2SiO3+4HCl . 

The  formation  and  properties  of  silicon  tetrafluoride,  SiF4,  and 
of  hydrofluosilicic  acid  and  its  salts  have  already  been  discussed 
(270-272). 

By  the  action  of  HC1  gas  on  heated  silicon,  silico-chloroform, 
SiHCl3,  is  obtained  as  a  low-boiling  colorless  liquid.  This  is  the 
analogue  of  chloroform,  CHC13  (644),  but  unlike  the  latter  it  is 
readily  decomposed  by  water.  Silicon  tetrachloride  vapor 
reacts  with  heated  silicon  as  follows: 

3SiCl4-f  Si-»2Si2Cl6 . 


520  Introduction  to  General  Chemistry 

The  product,  silicon  hexachloride,  is  a  liquid  boiling  at  148°, 
analogous  to  carbon  hexachloride,  C2C16,  boiling  at  187°.  There 
is  little  doubt  that  the  structure  of  these  substances  is  that 
represented  by  the  following  formulae: 

Cl   Cl  Cl    Cl 

Cl— C— C— Cl        Cl— Si— Si— Cl 

I      i  i      I 

Cl   Cl  Cl    Cl 

Much  more  complicated  silicon  compounds  are  known,  in 
some  of  which  long  chains  (664)  of  silicon  atoms  occur.  The 
resemblance  between  compounds  of  silicon  and  carbon  extends 
to  a  great  variety  of  compounds.  By  way  of  further  illustration 
the  following  pairs  of  formulae  of  related  substances  are  shown: 

CH3  CH3 

H3C— C— CH3  H3C-  Si— CH3 

I  I 

CH3  CH3 

Tetramethyl-M  ethane  Tetramethyl-Silico-M  ethane 

OCOH  OSiOH 

OCOH  OSiOH 

Oxalic  Acid  (665)  Silico-Oxalic  Acid 

OCOH  OSiOH 

HOCOH  HOSiOH 

OCOH  OSiOH 

Mesoxalic  Acid  Silico-Mesoxalic  Acid 

806.  Silica  or  Silicon  Dioxide,  SiO2. — Silica  is  one  of  the 
most  abundant  of  all  minerals.  It  occurs  in  many  forms  both 
crystalline  and  amorphous.  In  a  pure  state  it  forms  colorless, 
transparent,  hexagonal  crystals  called  quartz  or  rock  crystal. 
Amethyst  is  a  variety  of  quartz  colored  by  manganese.  The 
commonest  kinds  of  sand  are  usually  largely  or  wholly  quartz. 
Sandstone  is  essentially  quartz,  red  varieties  being  colored 


Some  Additional  Elements  and  Their  Compounds         521 

with  ferric  oxide.  A  massive  form  of  silica  is  called  quartzite. 
Amorphous  (581)  varieties  constitute  agate  and  onyx.  Flint 
and  opal  are  hydrated  varieties  of  silica. 

The  various  forms  of  silica  are  extremely  resistant  to  chemi- 
cal and  physical  changes  by  reason  of  the  great  hardness,  high 
melting-point,  and  chemical  inertness  of  this  substance.  These 
properties,  together  with  the  abundance  and  cheapness  of  various 
forms  of  silica,  cause  them  to  be  used  in  enormous  amounts  for 
a  great  variety  of  purposes.  As  an  abrasive  silica  finds  use  in 
grindstones,  whetstones,  sandpaper,  and  polishing  material,  the 
latter  containing  silica  powder,  tripoli,  or  infusorial  earth  (732). 
Mortar  and  concrete  always  contain  quartz  sand.  Fire  bricks  for 
furnace  linings  are  often  made  of  silica.  Various  kinds  of  glass, 
including  that  made  from  pure  quartz,  will  be  considered  later 

(808). 

807.  Silicic  Acid  and  Silicates. — Silica  reacts  with  sodium 
carbonate  (soda  ash)  at  a  red  heat  to  form  a  viscous  fluid  con- 
sisting of  sodium  silicate,  Na2SiO3.  During  the  fusion  carbon 
dioxide  is  evolved: 

SiO2+Na2CO3->Na2SiO3+C02 . 

On  cooling,  sodium  silicate  solidifies  to  a  glassy  solid  easily 
soluble  in  water.  By  adding  hydrochloric  acid  to  a  solution  of 
sodium  silicate  we  get  silicic  acid,  which  under  certain  conditions 
appears  as  a  gelatinous  precipitate  (a  gel,  761).  The  com- 
position of  this  silicic  acid  is  approximately  represented  by  the 
formula  H4SiO4.  Its  formation  may  be  represented  thus: 

Na2Si03+H2O+2HCl->H4SiO4-f  2NaCl . 

Silicic  acid  is  an  extremely  weak  acid,  but  it  will  dissolve  in 
sodium  hydroxide  solution  to  form  a  solution  of  sodium  silicate. 
The  solution  is  strongly  alkaline  by  reason  of  hydrolysis. 

Silicic  acid  gel  gradually  loses  water  when  heated,  and  in 
so  doing  it  doubtless  gives  a  series  of  derived  acids  and  finally 
silica.  In  this  respect  it  resembles  phosphoric  (590)  and  boric 
acids  (802).  Theoretically  we  have  the  following  possibilities 


522  Introduction  to  General  Chemistry 

as  one  or  two  molecules  of  H4Si04  lose  water.     (In  the  following 
expressions  the  water  lost  is  omitted  for  the  sake  of  brevity.) 

H4SiO4^H2SiO3->SiO2 
2H4SiO4->H6Si2O7->H4Si2O6^H2Si2O5->2SiO2. 

Still  more  complex  acids  would  result  from  the  dehydration 
of  three  or  more  molecules  of  silicic  acid.  Although  none  of 
these  acids  has  been  made  in  pure  form,  salts  of  several  of  them 
are  found  as  minerals.  Silicon  dioxide  is  of  course  the  anhydride 
(313)  of  all  the  possible  derivatives  of  silicic  acid,  and  the  formulae 
of  all  the  latter  may  be  written  in  the  general  form  (H2O)n(Si02)w, 
where  n  and  m  are  integers.  The  formulae  of  the  salts  may  also 
be  written  similarly,  as  illustrated  in  Table  XLIII,  which  shows 
the  composition  of  several  minerals  and  the  acids  from  which 
they  are  derived. 

TABLE  XLIII 


H4SiO4  

Zn2SiO4 

2ZnO  •  SiO2 

Willemite 

H2SiO3  

CaSiO3 

CaOSiO2 

Wollastonite 

H4SiO4  

2H2KAl3(SiO4)3 

K2O  •  3A12O3  •  6SiO2  •  2H2O 

Mica 

H2SiO3  

H2Mg3(SiO3)4 

3MgO-4Si(VH2O 

Talc 

H4Si308  

2KAlSi3O8 

K2O-Al2O3-6SiO2 

Orthoclase 

808.  Glass. — There  are  a  great  many  kinds  of  glass,  nearly 
all  of  which  are  silicates,  although  borates  are  also  common 
ingredients.  The  basic  constituent  elements  include  sodium 
potassium,  calcium,  aluminum,  zinc,  lead,  and  other  metals  less 
commonly  used.  Water  glass  having  approximately  the  com- 
position Na2Si4O9  is  made  by  dissolving  infusorial  earth  (732)  in 
hot  sodium  hydroxide  solution.  It  comes  on  the  market  as  a 
heavy  syrup  which  gives  on  evaporation  a  glasslike  solid  easily 
soluble  in  water.  It  is  used  as  a  component  of  laundry  soaps, 
as  a  cement,  as  an  oil-proof  glaze  for  lard  barrels,  and  for  a 
number  of  other  purposes. 

Glass  made  from  the  silicates  of  sodium  and  calcium  is  used 
for  window  panes,  bottles,  etc.  The  ordinary  soft  glass  tubing 
of  the  laboratory  is  also  made  of  this  kind  of  glass.  Hard  glass, 
which  contains  potassium  in  place  of  the  sodium  in  soft  glass, 


Some  Additional  Elements  and  Their  Compounds        523 

fuses  at  a  higher  temperature  than  the  latter.  Flint  glass,  a 
potassium-lead  silicate,  has  a  high  index  of  refraction  and  on 
this  account  possesses  great  brilliancy.  It  is  used  for  cut-glass 
ware  and  lenses. 

Glassware  for  chemical  use,  beakers,  flasks,  etc.,  should 
possess  several  special  properties.  It  must  be  practically 
insoluble  in  water,  acids,  and  alkalies,  must  not  crack  when 
heated,  and  must  be  thick  enough  to  withstand  ordinary  handling 
without  breaking.  The  cracking  of  glass  when  heated  is  the 
result  of  rapid  local  expansion.  The  thinner  the  glass  the  less 
liable  it  is  to  crack  when  heated;  but  if  ordinary  glass  is  made 
thin  enough  to  withstand  heating  in  the  usual  way  it  is  very 
fragile.  Before  the  war  nearly  all  chemical  glassware  came  from 
Germany  and  Bohemia.  There  were  two  principal  kinds,  typi- 
cally represented  by  Kavalier  glass  and,  Jena  glass,  the  latter 
being  much  better  in  quality  than  the  former.  The  composition 
of  each  of  these  is  shown  in  Table  XLIV. 


TABLE  XLIV 
PERCENTAGE  COMPOSITION  OF  GLASS  FOR  LABORATORY  WARE* 


Kavalier 

Jena 

Pyrex 

Nonsol 

A12O3-  • 

o.  14 

4.  20 

2.00 

2.50 

Fe2O3 

O    2< 

O.  2< 

o.  23 

ZnO         .    . 

IO    QO 

7.80 

CaO  
MgO.  . 

8.70 

0.63 
O.  21 

0.29 
O.06 

0.79 

3.40 

K20  
Na2O  

7.90 
7.  10 

0-37 
7.50 

O.  2O 
4.40 

0.30 
10.90 

SiO2  

75.90 

64.  70 

80.50 

67.30 

B2O, 

10  90 

II.80 

6.  20 

As^O, 

O    14. 

o  70 

*  Walker  &  Smith,  U.S.  Bureau  of  Standards,  Bulletin  107, 1918. 

During  the  war  several  American  manufacturers  succeeded 
in  producing  first-class  chemical  glass.  The  composition  of  two 
sorts,  Pyrex  and  Nonsol,  is  shown  in  Table  XLIV.  Pyrex  glass 
has  so  small  a  coefficient  of  expansion  that  beakers  and  flasks 
can  be  made  with  thick  walls  and  still  be  far  less  liable  to  crack 
with  change  of  temperature  than  the  best  German  glass.  Several 


524  Introduction  to  General  Chemistry 

other  sorts  of  American  chemical  glass  are  also  of  excellent 
quality,  so  that  today  we  have  better  domestic  ware  than  was 
heretofore  imported. 

Quartz  glass  is  made  from  pure  silicon  dioxide  (806)  melted 
in  an  electric  furnace.  This  glass  resembles  ordinary  glass  in 
appearance  but  differs  from  the  latter  in  having  an  extremely 
high  melting-point  and  an  extremely  small  coefficient  of  expan- 
sion. White-hot  quartz  will  not  crack  when  plunged  into  water. 
This  glass  finds  important  uses  in  both  scientific  and  technical 
chemistry. 

809.  Tin  and  Its  Compounds. — The  metal  tin  (Sn=  119)  was 
known  from  very  early  times,  its  Latin  name  being  stannum. 
It  is  widely  used  in  making  a  number  of  alloys,  such  as  bronze 
and  bell  metal  (tin  and  copper),  solder  (tin  and  lead),  and  tin 
amalgam  (tin  and  mercury  used  in  making  mirror  backs) .  Sheet 
iron  coated  with  tin  and  known  as  tin  plate  is  made  in  immense 
quantities  for  the  manufacture  of  common  tinware.  Tin  in  the 
form  of  thin  sheets  is  tin  foil,  but  the  common  grade  of  this 
article  is  usually  largely  lead. 

Tin  is  a  rather  soft  metal  of  low  melting-point,  232°.  It  is 
permanent  in  air  and  in  water,  and,  as  shown  by  its  position  in 
the  displacement  series  (Table  XIX,  492),  it  is  a  rather  inactive 
element.  It  dissolves  somewhat  slowly  in  hydrochloric  acid, 
forming  stannous  chloride,  SnCl2, 

Sn+2HCl->SnCl2+H2. 

The  hydrate  SnCl2  •  2H2O,  easily  soluble  in  water,  is  known  in 
commerce  as  tin  salt  and  is  used  in  dyeing.  It  is  frequently  used 
in  the  laboratory  as  a  reducing  agent.  Stannous  sulfate,  SnSO4, 
and  nitrate,  Sn(N03)2,  can  also  be  made.  The  action  of  chlorine 
gas  on  melted  tin  gives  stannic  chloride,  SnCl4, 

Sn+2d2->SnCl4 . 

The  product  is  a  colorless  liquid  boiling  at  114°,  which  forms 
with  water  the  crystalline  hydrates  SnCl4  •  5H2O  and  SnCl4  •  8H2O. 
These  salts  are  also  used  as  mordants  in  dyeing,  a  mordant  being 
a  substance  which  by  combining  with  the  fiber  fixes  the  dye  upon 
the  latter.  Numerous  other  stannic  salts  are  known. 


Some  Additional  Elements  and  Their  Compounds         525 

Thus  it  will  be  seen  that  tin  forms  two  series  of  salts:  the 
stannous,  in  which  the  metal  is  bivalent,  and  the  stannic,  in 
which  it  is  quadrivalent.  Stannous  salts  form  stannous  ions, 
Sn++,  which  by  oxidation  give  stannic  ions,  Sn4+.  The  latter 
by  reduction  give  the  former.  Stannous  solutions  are  strong 
reducing  agents. 

Tin  forms  two  oxides,  stannous,  SnO,  and  stannic,  Sn02. 
The  corresponding  hydroxides  are  precipitated  when  a  stannous 
or  a  stannic  solution  is  treated  with  the  equivalent  amount  of 
sodium  hydroxide.  Each  precipitate  is  soluble  in  an  excess  of 
the  alkali,  thus  showing  that  the  hydroxides  of  tin  have  acidic 
as  well  as  basic  properties.  Like  aluminum  hydroxide  they  are 
amphoteric  (177).  Stannous  hydroxide,  Sn(OH)2,  yields  sodium 
stannite,  Na2SnO2;  while  stannic  hydroxide,  also  called  stannic 
acid,  gives  sodium  stannate,  Na2Sn03.  A  solution  of  sodium 
stannite  is  a  very  powerful  reducing  agent,  since  it  is  readily 
oxidized  to  sodium  stannate.  An  isomer  (647)  of  stannic  acid, 
called  metastannic  acid,  is  formed  by  the  action  of  hot  nitric 
acid  on  tin.  This  is  a  white  solid,  insoluble  in  water  and  ni- 
tric acid,  and  differs  greatly  from  the  other  form.  Metastannic 
acid  forms  complex  salts  such  as  Na2Sn5Ou'4H20.  Hydrogen 
sulfide  precipitates  brown  stannous  sulfide  from  solutions  of 
stannous  chloride  and  yellow  stannic  sulfide  from  solutions  of 
stannic  chloride.  Both  precipitates  are  soluble  in  concentrated 
hydrochloric  acid.  Stannic  sulfide  is  dissolved  by  ammonium 
sulfide,  owing  to  the  formation  of  ammonium  sulfo-stannate. 
Stannous  sulfide  does  not  dissolve  appreciably  in  pure  ammonium 
sulfide  but  does  so  in  yellow  ammonium  sulfide  (607),  forming 
also  ammonium  sulfo-stannate: 

SnS2+(NH4)2S  -^(NH4)2SnS3 , 
SnS  +  (NH4)2S2->(NH4)2SnS3 . 

810.  Arsenic. — In  its  chemistry  arsenic  (As  =  75)  is  much 
more  like  the  non-metallic  than  the  metallic  elements.  In 
many  of  its  compounds  it  bears  a  close  resemblance  to  phos- 
phorus. Arsenic  is  an  abundant  element,  occurring  as  a  mineral 
both  free  and  in  combination  with  oxygen  and  sulfur,  and  with 


526  Introduction  to  General  Chemistry 

iron,  copper,  and  many  other  metals.  The  trioxide  As2O3, 
commonly  called 'white  arsenic,  is  obtained  in  large  amounts  as 
a  by-product  of  the  smelting  of  ores  of  copper  and  some  other 
metals. 

The  free  element  is  easily  prepared  by  heating  the  trioxide 
with  carbon;  the  vapor  of  arsenic  so  formed  condenses  on  cooling 
to  nearly  black  crystals.  The  latter  are  semimetallic  in  appear- 
ance and,  like  metals,  conduct  electricity  well.  Arsenic  is  not 
oxidized  by  air  at  ordinary  temperatures  but  burns  when  strongly 
heated,  forming  the  trioxide.  When  heated  in  a  stream  of 
chlorine,  arsenic  forms  a  trichloride,  AsCl3,  a  colorless  liquid 
boiling  at  130°.  This  hydrolyzes  almost  completely  when 
treated  with  much  water,  thus  showing  that  the  corresponding 
hydroxide,  As(OH)3,  is  a  very  weak  base. 

On  the  other  hand  the  hydroxide  and  also  the  trioxide  dis- 
solve readily  in  dilute  alkalies  to  form  salts  in  which  As(OH)3 
plays  the  part  of  an  acid,  called  arsenious  acid.  Two  series  of 
salts  are  known:  one  derived  from  H3AsO3,  the  other  from 
HAs02.  Arsenious  acid,  like  phosphorous  acid,  H3PO3  (588), 
is  a  good  reducing  agent.  It  is  oxidized  by  halogens  or  nitric 
acid  to  form  arsenic  acid,  H3As04: 

H3As03+H20+Cl3->H3AsO4+2HCl . 

Arsenic  acid,  which  is  derived  from  the  pentoxide  As2Os,  is  a 
very  soluble,  moderately  strong  acid  closely  resembling  phos- 
phoric acid  (590).  Like  the  latter,  it  loses  water  when  heated, 
giving  first  pyroarsenic  acid,  H4As2O7,  and  then  metarsenic  acid, 
HAs03.  Arsenic  acid  resembles  phosphoric  acid  in  giving  with 
ammonium  molybdate  solution  a  yellow  precipitate  of  am- 
monium arseno-molybdate,  insoluble  in  nitric  acid  (597). 

The  numerous  salts  of  arsenic  acid  closely  resemble  the 
corresponding  salts  of  phosphoric  acid.  In  fact  Mitscher- 
lich  in  the  year  1819  discovered  that  corresponding  salts,  e.g., 
H2KP04  and  H2KAsO4,  were  identical  in  crystalline  form.  Such 
compounds  were  said  to  be  isomorphous.  Later  Mitscherlich 
declared  that  if  two  substances  are  isomorphous  they  must  have 
the  same  numbers  of  atoms  similarly  arranged  in  the  molecule. 


Some  Additional  Elements  and  Their  Compounds        527 

This  principle  enabled  him  to  discover  the  correct  formulae  of 
many  compounds  of  little-known  elements  at  a  time  when  much 
uncertainty  prevailed  regarding  valence  and  atomic  weights. 

The  action  of  hydrogen  sulfide  on  a  solution  of  arsenious 
acid  acidified  with  HC1  gives  a  precipitate  of  yellow  arsenic 
trisulfide,  As2S3  (608).  Acidified  solutions  of  arsenic  acid  or 
its  salts  give  under  similar  conditions  arsenic  pentasulfide, 
As2S5,  along  with  some  trisulfide  and  sulfur,  according  to  experi- 
mental conditions.  These  sulfides  are  insoluble  in  water  and 
hydrochloric  acid  but  are  easily  soluble  in  concentrated  nitric 
acid,  which  changes  the  substances  into  sulfuric  and  arsenic 
acids.  The  sulfides  also  dissolve  in  ammonium  sulfide  to  form 
soluble  sulfarsenite,  (NH4)3AsS3,  and  sulfarsenate,  (NH4)3AsS4. 

Arsenic  and  hydrogen  form  a  gaseous  compound  called 
arsine,  AsH3.  This  is  produced  along  with  hydrogen  when  any 
soluble  arsenic  compound  is  added  to  a  mixture  of  zinc  and 
hydrochloric  acid.  It  is  an  extremely  poisonous  gas  and  must 
be  handled  with  great  caution.  Arsine  is  readily  decomposed 
into  its  elements  by  heat,  so  that  if  it  is  passed  through  a  glass 
tube  heated  locally  a  dark  mirror-like  deposit  of  arsenic  is 
formed  in  the  tube.  Marsh's  test  for  arsenic  is  based  on  these 
reactions. 

Most  arsenic  compounds  are  poisonous;  one-fifth  of  a  gram 
of  the  trioxide  usually  proves  fatal.  A  good  antidote  consists 
of  freshly  prepared  ferric  hydroxide,  which  combines  readily 
with  arsenious  acid  to  form  an  insoluble  and  therefore  little- 
poisonous  compound. 

811.  Antimony. — Antimony  (Sb  =  120)  is  classed  as  a  metallic 
element  because  of  the  decidedly  metallic  properties  of  the  free 
element  and  its  ability  to  form  salts  with  acids  such  as  sul- 
furic and  nitric.  Nevertheless  its  hydroxides  are  also  acidic 
and  give  rise  to  salts  with  bases,  just  as  do  the  hydroxides  of 
arsenic.  Antimony  is  a  widely  distributed  and  rather  common 
element  occurring  usually  in  combination  with  oxygen,  sulfur, 
or  metallic  elements.  China  is  the  largest  single  source  of  the 
element.  The  metal  forms  brittle,  silver-white  crystals,  melting 
at  630°.  It  forms  useful  alloys  with  many  other  metals.  Type 


528  Introduction  to  General  Chemistry 

metal,  an  alloy  with  lead  and  bismuth,  expands  at  the  moment 
of  solidifying  in  the  mold  and  thus  produces  type  with  sharp, 
clear-cut  edges.  Britannia,  made  from  tin,  antimony,  and  copper, 
is  used  as  a  cheap  substitute  for  silver  plate.  Babbitt  metal, 
used  for  antifriction  bearings,  consists  of  antimony  and  tin  with 
someJead  and  copper. 

Antimony  unites  readily  with  chlorine  to  form  the  trichloride 
SbCl3,  white  crystals  melting  at  73°  and  boiling  at  223°.  The 
chloride  is  to  be  considered  the  salt  of  the  weak  base  Sb(OH)3; 
it  is  soluble  in  aqueous  hydrochloric  acid  but  is  hydrolyzed  by 
pure  water.  With  a  small  proportion  of  water  a  white  precipitate 
of  sparingly  soluble  antimony  oxychloride,  SbOCl,  is  formed, 

SbCl3+H2O->SbOCl-f  2HC1 . 

This  reaction  is  reversible,  SbOCl  dissolving  readily  in  concen- 
trated HC1.  The  univalent  radical  SbO  present  in  several  salts 
is  called  antimonyl;  therefore  SbOCl  is  also  called  antimonyl 
chloride. 

Antimony  nitrate,  Sb(NO3)3,  and  antimony  sulfate,  Sb2(SO4)3, 
are  typical  salts  as  far  as  their  composition  is  concerned,  but  they 
are  more  or  less  completely  hydrolyzed  by  water,  depending 
on  the  proportion  and  temperature  of  the  latter.  The  double 
salts  potassium  antimonyl  tartrate,  K(SbO)C4H406  (see  665), 
and  sodium  antimony  fluoride,  NaSbF4,  are  easily  soluble  salts 
which  are  not  hydrolytically  decomposed  by  water.  Both  of 
these  salts  are  important  mordants  (809),  and  the  former,  com- 
monly known  as  tartar  emetic,  is  also  used  in  medicine.  In 
both  cases  stable  complex  ions  (538)  are  formed;  these  are 
(SbO)C4H406-  and  SbF4~  respectively. 

Antimony  forms  three  oxides:  Sb2O3,  Sb204,  and  Sb2Os. 
The  first  and  second  can  be  made  by  direct  union  of  antimony 
with  oxygen,  the  second  being  the  one  stable  in  air  at  a  red  heat. 
The  third,  made  indirectly,  loses  oxygen  when  strongly  heated, 
giving  Sb204.  The  trioxide  corresponds  to  a  hydroxide,  Sb (OH) 3, 
or  H3Sb03,  which  by  reason  of  its  faintly  acidic  nature  is  known 
as  antimonous  acid.  The  latter  unites  with  bases  to  form  salts 
such  as  Na3SbO3  and  NaSbO2  •  3H2O.  Antimony  pentoxide  is 


Some  Additional  Elements  and  Their  Compounds         529 

the  anhydride  of  antimonic  acid,  H3SbO4,  a  white  powder  spar- 
ingly soluble  in  water.  This  weak  acid  loses  water  when  heated, 
giving  in  turn  pyroantimonic  acid,  H4Sb2O7,  metantimonic  acid, 
HSbO3,  and  finally  Sb20s.  Potassium  hydrogen  pyroantimonate, 
K2H2Sb207,  is  one  of  the  best-known  salts  of  these  acids.  Its 
solution  gives  with  sodium  salts  a  precipitate  of  sodium  hydrogen 
pyroantimonate,  Na2H2Sb2O7  •  6H20. 

Acid  solutions  of  antimony  salts  give  with  hydrogen  sulfide 
brick-red  precipitates  of  antimony  trisulfide,  Sb2S3,  or  penta- 
sulfide,  Sb2Ss.  These  sulfides  are  almost  insoluble  in  water  and 
cold  dilute  acids  but  dissolve  easily  in  potassium  hydroxide  by 
reason  of  reactions  like  the  following: 

2Sb2S3+4KOH->3KSbS2+KSbO2+2H20 . 

The  soluble  salt  KSbS2  is  called  potassium  sulfantimonite. 

Antimony  trisulfide  also  dissolves  in  yellow  ammonium 
sulfide  to  form  soluble  ammonium  sulfantimonate,  (NH4)3SbS4. 

Antimony  forms  a  gaseous  hydride,  SbH3,  called  stibine. 
This  gas  resembles  arsine  closely  and  is  also  decomposed  by 
heat,  with  the  formation  of  a  black  deposit  of  free  antimony. 

812.  Bismuth. — Bismuth  (Bi  =  2o8)  is  strictly  metallic  in 
properties,  in  distinction  from  arsenic,  which  is  non-metallic, 
and  antimony,  which  stands  midway  between  the  two  in  this 
respect.  In  its  ordinary  compounds  bismuth  is  trivalent,  but 
products  probably  pentavalent  also  exist.  The  element  is  much 
scarcer  than  either  arsenic  or  antimony  and  commands  a  much 
higher  price.  It  occurs  free  (native)  and  as  oxide  or  sulfide, 
often  admixed  with  ores  of  copper  or  lead.  Bolivia  is  the 
chief  source  of  bismuth.  A  considerable  amount  of  it  is  obtained 
in  the  United  States  as  a  by-product  in  the  electrolytic  refining 
of  lead.  The  metal  forms  rather- hard,  brittle  crystals,  silver- 
white  of  a  reddish  tint.  The  melting-point  is  269°.  The  metal 
is  used  in  the  making  of  easily  fusible  alloys.  Thus  Rose's  metal, 
consisting  of  tin  i,  lead  i,  and  bismuth  2  parts,  melts  at  94°; 
Wood's  metal,  tin  i,  lead  2,  bismuth  4,  and  cadmium  i,  melts 
at  61°.  Fusible  alloys  are  used  for  fire-sprinkler  nozzles,  electric 
fuse  wires,  etc. 


530  Introduction  to  General  Chemistry 

Bismuth  forms  with  acids  salts  in  which  this  element  acts 
as  a  trivalent  positive  ion.  Bismuth  trichloride,  BiCl3,  meltint 
at  227°  and  boiling  at  428°,  is  made  by  the  direct  union  of  chlorine 
with  the  metal,  or  by  the  action  of  aqua  regia  (562)  on  the  lat- 
ter. The  salt  is  soluble  in  water,  but  when  diluted  the  salg 
hydrolyzes  and  forms  a  white  precipitate  of  bismuth  oxychloride, 
BiOCl  (bismuthyl  chloride  or  basic  bismuth  chloride).  Bismuth 
nitrate,  Bi(NO3)3,  results  from  the  action  of  nitric  acid  on  the 
metal.  It  is  soluble  in  water  or  dilute  nitric  acid  but  is  hydro- 
lyzed  by  much  water,  forming  the  basic  nitrate  BiON03  (also 
called  bismuth  subnitrate).  This  product  is  a  nearly  insoluble 
white  powder  used  both  internally  and  externally  in  medicine. 
Quite  large  doses  are  easily  tolerated.  Bismuth  sulfide,  Bi2S3, 
is  formed  as  a  brown  precipitate  when  acid  solutions  of  bismuth 
salts  are  treated  with  hydrogen  sulfide.  Sodium  hydroxide  or 
ammonia  precipitates  white  bismuth  hydroxide,  Bi(OH)3,  from 
a  bismuth  salt  solution.  Bismuth  does  not  form  a  hydride. 

813.  Molybdenum. — Molybdenum  (Mo  =  96)  is  found  in 
nature  chiefly  as  molybdenite,  MoS2,  a  mineral  closely  resembling 
graphite  (630)  in  appearance.  When  the  sulfide  is  heated  in 
air  (roasted)  it  is  oxidized  to  the  trioxide  Mo03.  The  latter 
heated  with  hydrogen  or  carbon  gives  the  metallic  element. 
Molybdenum  is  a  hard  but  malleable  metal  of  very  high  melting- 
point  (over  2400°) .  It  is  not  oxidized  by  air  at  ordinary  tempera- 
tures. The  pure  metal  is  not  commercially  important,  but  an 
alloy  with  iron,  ferro-molybdenum,  is  made  on  a  rather  large 
scale.  This  alloy  finds  use  in  the  making  of  certain  kinds  of  high- 
grade  tool  steel. 

The  trioxide  Mo03  is  the  most  important  of  the  several 
oxides.  It  is  the  anhydride  of  molybdic  acid,  H2Mo04  •  H20, 
which  consists  of  yellow  crystals  difficultly  soluble  in  water.  The 
hydrate  loses  water  when  heated  gently,  giving  white  H2Mo04? 
Sodium  molybdate,  Na2Mo04«ioH2O,  resembles  Glauber's  salt. 
Na2S04'ioH20,  in  appearance.  Molybdic  acid  resembles  silicic 
acid  in  forming  salts  of  complex  formulae  (807).  Thus  for 
example  a  sodium  salt,  Na6Mo7024  •  22H20,  and  an  ammonium 
salt,  (NH4)6Mo7024-4H20,  are  known.  The  latter  salt  dis- 


Some  Additional  Elements  and  Their  Compounds        531 

solved  in  an  excess  of  nitric  acid  gives  the  so-called  ammonium 
molybdate  solution,  extensively  used  in  the  detection  and  quanti- 
tative analysis  of  phosphates  (597).  With  this  reagent  a  nitric 
acid  solution  of  a  phosphate  or  phosphoric  acid  gives  quantita- 
tively a  yellow  precipitate  of  ammonium  phospho-molybdate, 
(NH4)3P04'iiMoO3-6H20.  Arsenic  acid  solutions  give  the 
corresponding  arsenic  compound  under  similar  conditions  (810). 

814.  Tungsten. — The  element  tungsten  (W=i84)  is  of  con- 
siderable practical  importance.  In  the  United'States  it  is  found 
chiefly  in  Colorado.  Its  principal  ore  is  Wolframite,  a  tungstate 
of  iron  and  manganese.  The  separate  tungstates  FeWO4  and 
MnWO4  also  occur  as  minerals.  Metallic  tungsten  is  made  in 
very  pure  form  by  the  reduction  of  its  oxide,  W03,  by  hydrogen 
at  a  high  temperature.  By  reason  of  the  extremely  high 
melting-point  of  the  metal  the  product  is  obtained  as  a  powder. 
To  get  this  into  compact  malleable  and  ductile  form  the  powder 
is  made  into  bars  by  powerful  compression;  these  are  then  again 
heated  electrically  and  the  granules  caused  to  melt  together. 
Pure  metallic  tungsten,  first  made  by  Cooledge  in  1911,  is  a 
hard,  silver-white,  and  very  ductile  metal.  Its  most  important 
use  is  for  the  manufacture  of  the  filaments  of  incandescent 
electric  lamps.  Its  great  value  for  this  purpose  depends  on  the 
fact  that  its  melting-point,  3267°,  is  higher  than  that  of  any  other 
metal.  It  may  be  added  that  the  efficiency  of  an  electric  lamp 
(ratio  of  light  energy  to  electrical  energy)  increases  greatly  with 
the  temperature  to  which  the  filament  can  be  heated. 

Ferro-tungsten  (70  per  cent  W,  30  per  cent  Fe),  made  on  a 
large  scale  by  an  electric-furnace  process,  is  used  in  the  manu- 
facture of  special  steels.  Tungsten  forms  two  oxides,  W02  and 
W03,  the  latter  being  far  more  important.  The  trioxide  is  the 
anhydride  of  tungstic  acid,  H2W04,  a  yellow  powder  insoluble  in 
water.  This  acid  forms  simple  tungstates  such  as  Na2W04  •  2H20, 
K2W04,  CaWO4,  PbWO4,  and  the  tungstates  of  iron  and  man- 
ganese already  mentioned.  Calcium  tungstate,  CaW04,  gives 
a  fluorescent  light  with  X-rays  and  finds  important  use  for  the 
manufacture  of  fluoroscopic  screens.  Sodium  tungstate  is  used 
to  render  cotton  cloth  slow-burning  (" fireproof "). 


532  Introduction  to  General  Chemistry 

815.  Selenium  and  Tellurium. — The  elements  selenium 
(Se  =  79)  and  tellurium  (Te=i27-5)  are  more  closely  related  to 
sulfur  than  to  any  others.  The  first,  selenium,  is  essentially 
non-metallic  in  its  behavior.  The  second,  tellurium,  in  free 
form  is  a  silver- white  crystalline  metal  which,  in  addition  to 
forming  some  salts  in  which  it  is  the  basic  ion,  also  forms  two 
well-characterized  acids.  These  elements  are  not  very  common. 
Practically  they  are  obtained  as  by-products  in  the  refining  of  cop- 
per. No  important  use  has  been  found  for  tellurium.  Selenium 
has  two  uses,  one  in  the  manufacture  of  red  glass,  the  other  in  the 
construction  of  light-sensitive  electrical  apparatus.  The  electri- 
cal conductivity  of  selenium  varies  greatly  with  the  intensity  of 
its  illumination,  so  that  by  ingenious  use  of  this  principle  pic- 
tures, etc.,  have  been  transmitted  long  distances  by  wire.  Lights 
on  beacons  and  buoys  may  be  turned  on  at  dark  by  the  use  of 
selenium  cells.  Selenium  and  tellurium  form  compounds  analo- 
gous to  H2S  by  methods  similar  to  those  that  yield  the  latter  gas. 

Hydrogen  selenide,  H2Se,  is  an  ill-smelling  gas,  dangerously 
irritating  to  mucous  membranes  and  temporarily  paralyzing 
the  sense  of  smell.  Hydrogen  telluride,  H2Te,  is  a  similar  gas. 
Selenium  and  tellurium  both  burn  in  air  or  oxygen  when  ignited 
and  form  white  solid  dioxides,  Se02  and  TeO2.  These  oxides 
both  unite  with  water  to  form  selenous  acid  and  tellurous  acid 
respectively : 

Se02+H20->H2Se03 , 

TeO2+H2O->H2Te03 . 

These  acids  are  the  analogues' of  sulfurous  acid,  H2S03  (611), 
but  are  more  stable  than  the  latter  in  that  both  are  white 
crystalline  solids.  They  form  typical  salts  with  bases;  not 
only  normal  salts  like  K2Se03  and  K2Te03,  but  also  acid  salts 
like  KHSe03  and  KHTeO3. 

Unlike  sulfurous  acid,  these  acids  are  oxidizing  agents,  since 
they  are  reduced  to  selenium  and  tellurium  respectively  by 
reducing  agents,  as  illustrated  in  the  following  equation: 

H2Se03+2S02+H20->Se-}-2H2S04 . 

On  the  other  hand,  powerful  oxidizing  agents  convert  selenous 
and  tellurous  acids  into  selenic  acid,  H2Se04,  and  telluric  acid, 


Some  Additional  Elements  and  Their  Compounds        533 

H2TeO4,  respectively.  These  acids  form  white  crystals  and  are 
analogous  in  composition  to  sulfuric  acid.  Like  the  latter,  they 
form  both  neutral  and  acid  salts,  e.g.,  potassium  acid  tellurate, 
KHTe04.  The  selenates  and  tellurates  are  oxidizing  agents 
tending  to  pass  into  selenites  and  tellurites.  Thus,  for  example, 
potassium  tellurate,  when  strongly  heated,  gives  the  tellurite 
and  oxygen,  or  when  heated  with  concentrated  hydrochloric 
acid  it  gives  potassium  tellurite  and  chlorine.  Telluric  acid  is 
decomposed  by  heat  into  its  anhydride,  tellurium  trioxide,  TeO3, 
and  water. 

816.  Cobalt  and  Nickel.— The  metals  cobalt  (00  =  58.97) 
and  nickel  (Ni=58.68)  resemble  iron  in  both  their  physical 
and  their  chemical  properties  but  are  less  active  chemically, 
corresponding  to  positions  below  iron  in  the  displacement  series 
(492).  Nickel  is  by  far  the  more  important  of  the  two  com- 
mercially, although  cobalt  commands  the  higher  price.  The 
two  metals  usually  occur  together,  nickel  being  the  more  abun- 
dant. They  are  commonly  found  with  iron,  copper,  or  silver  as 
sulfides  or  arsenides.  Large  deposits  of  nickel  and  cobalt  occur 
in  the  province  of  Ontario,  Canada. 

A  unique  method  of  extracting  nickel  from  its  ores  is  based 
on  the  fact  that  carbon  jnonoxide,  CO  (632),  unites  with  finely 
divided  metallic  nickel  at  80°  to  form  a  compound,  Ni(CO)4, 
nickel  carbonyl,  which  is  gaseous  above  44°.  The  ores,  after  they 
have  been  roasted  to  remove  sulfur  and  arsenic  and  to  convert 
the  metals  into  oxides,  are  first  extracted  with  sulfuric  acid  to 
remove  copper  and  are  then  reduced  to  metal  by  water-gas,  a 
mixture  of  H2  and  CO,  at  400°.  The  residue,  cooled  to  80°,  is 
exposed  to  a  stream  of  highly  compressed  CO,  and  the  volatile 
Ni(CO)4  which  passes  off  is  decomposed  at  200°  into  pure 
metallic  nickel  and  carbon  monoxide, 

Ni(CO)4->Ni-f  4CO . 

This  is  the  Mond  process.  Other  processes  of  refining  nickel  are 
also  in  use. 

Nickel  is  so  extensively  used  for  plating  copper,  brass,  and 
iron  that  its  appearance  and  permanence  in  the  air  are  known  to 
everyone.  It  is  harder  than  iron  and  melts  at  1450°.  The  metal 


534  Introduction  to  General  Chemistry 

finds  one  of  its  most  important  uses  in  the  manufacture  of 
nickel  steel,  of  which  it  may  constitute  several  per  cent.  Cobalt 
may  also  be  present  to  advantage  in  such  steel.  Nickel  steel  is 
very  hard  and  tough  and  is  largely  used  for  armor  plate. 

Our  five-cent  coins,  "nickels,"  are  an  alloy  of  25  per  cent 
nickel  and  75  per  cent  copper.  The  well-known  properties 
of  this  alloy  need  no  comment.  The  alloy  called  German 
silver  consists  of  nickel,  copper,  and  zinc.  Monel  metal  is 
Ni  68  per  cent,  Cu  30 . 5  per  cent,  and  Fe  i .  5  per  cent.  It 
resembles  the  alloy  of  our  nickel  coins  but  is  harder  and  more 
resistant  to  chemicals.  It  is  obtained  by  direct  smelting  of 
copper-nickel-iron  ores. 

Cobalt  resembles  nickel  in  appearance  and  properties.  It  is 
not  much  used  as  a  metal.  Its  oxide,  CoO,  is  used  in  making 
blue  glass,  and  since  no  other  means  is  known  of  making  a  good 
blue  glass  this  element  is  of  considerable  importance.  Blue 
glazes  for  porcelain  and  a  blue  pigment,  smalt,  are  also  made 
from  cobalt  oxide. 

817.  Salts  of  Cobalt  and  Nickel. — The  bivalent  oxides  and 
hydroxides  of  cobalt  and  nickel  are  basic  and  have  no  acidic 
properties.  They  form  salts  with  practically  all  acids.  In  its 
salts  nickel  is  always  and  cobalt  usually  bivalent,  although  a 
few  trivalent  salts  of  the  latter  element  are  known.  The  salts 
of  these  elements  correspond  closely  in  composition  to  those  of 
ferrous  iron.  Of  the  commoner  salts  of  (bivalent)  cobalt  and 
nickel  we  may  mention  the  fluorides,  chlorides,  bromides,  iodides, 
nitrates,  sulfates,  carbonates,  and  phosphates.  These  salts, 
with  the  exception  of  the  carbonates  and  phosphates,  are  readily 
soluble  in  water  and  as  a  rule  form  crystals  with  six  molecules  of 
water  to  one  of  the  salt,  e.g.,  Co(N03)2  •  6H2O  and  NiCl2  •  6H2O. 
Nickel  sulfate  forms  besides  the  hexahydrate  a  heptahydrate, 
NiS04  •  ;H2O,  isomorphous  (810)  with  MgSO4  •  7H2O  (144). 
Cobalt  sulfate  also  forms  a  hexahydrate  and  a  heptahydrate> 
the  latter  isomorphous  with  FeS04  •  7H2O  (173).  Many 
double  salts  (175)  are  also  known.  Ammonium  nickel  sul- 
fate, (NH4)2SO4  •  NiS04  •  6H20,  is  used  in  electroplating. 

The  hydrated  salts  of  cobalt  are  red  in  crystalline  form  and 
also  in  solution,  in  which  latter  Co++  ions  are  present.  Many 


Some  Additional  Elements  and  Their  Compounds         535 

anhydrous  cobalt  salts  are  blue.  If  one  uses  a  dilute  cobalt 
chloride  solution,  pale  red  in  color,  as  ink,  the  writing  will  be 
practically  invisible,  when  air  dry.  If  the  paper  is  heated  the 
writing  appears  in  easily  legible  blue  characters.  The  effect  of 
the  heating  is  to  drive  off  the  water  of  hydration.  A  writing- 
fluid  of  this  kind  is  known  as  sympathetic  ink.  Cobalt  forms  a 
large  number  of  compounds  called  cobalt  amines.  These  are 
complex  substances  containing  ammonia  and  trivalent  cobalt. 
Hexamine  cobaltic  chloride,  (NH3)6CoCl3,  formed  by  the 
action  of  ammonia  and  oxidizing  agents  on  a  solution  of  cobaltous 
chloride,  CoCl2,  is  a  typical  example.  Solutions  of  cobalt 
amines  must  contain  complex  ions  (538),  since  they  do  not  give 
the  ordinary  reactions  of  solutions  of  simple  cobalt  salts.  Nickel 
does  not  form  analogous  compounds. 

Hydrated  nickel  salts  are  green  and  form  green  solutions 
in  which  bivalent  Ni++  ions  are  present.  Simple  cobalt  and 
nickel  solutions  give  precipitates  with  alkaline  hydroxides, 
carbonates,  phosphates,  and  sulfides.  These  are  all  soluble  in 
ammonia  solution  by  reason  of  the  formation  of  complex  ions 
of  various  sorts. 

818.  Platinum. — Platinum  (Pt=i95.2)  occurs  only  in  the 
metallic  state,  often  as  an  alloy  with  the  rarer  metals  osmium 
and  indium.  It  comes  chiefly  from  the  Ural  Mountains  in 
Russia.  Platinum  is  almost  indispensable  in  chemical  labora- 
tories, where  it  is  used  as  crucibles,  dishes,  tubes,  etc.  Its 
value  for  such  uses  arises  from  the  fact  that  it  is  very  inactive 
and  has  a  high  melting-point.  It  is  not  attacked  by  any  of  the 
common  acids  singly,  although  it  is  slowly  converted  into  its 
chloride,  PtCl4,  by  aqua  regia  (562),  a  mixture  of  hydrochloric 
and  nitric  acids.  Platinum  is  an  active  catalytic  agent  for  many 
reactions,  such  as  the  union  of  hydrogen  and  oxygen  (303),  the 
union  of  sulfur  dioxide  and  oxygen  (617),  and  the  oxidation  of 
ammonia  to  nitric  acid  (570) .  It  finds  extensive  use  catalytically 
in  the  manufacture  of  sulfuric  and  nitric  acids.  Since  platinum 
has  in  recent  years  become  much  more  valuable  than  gold,  it  has 
gained  great  popularity  as  a  setting  for  jewels.  Its  use  in  jewelry 
was  restricted  and  later  prohibited  during  the  war  in  order  that 
the  metal  might  be  conserved  for  chemical  uses.  The  use  in 


Introduction  to  General  Chemistry 

jewelry  of  this  noble  metal,  so  necessary  for  chemistry  and  in- 
dustry, ought  permanently  to  be  prohibited.  The  compounds  of 
platinum  are  not  of  much  practical  importance.  Until  recently 
the  chloride  PtCl4,  a  very  soluble  yellow  salt,  was  used  in  the 
quantitative  determination  of  potassium.  Potassium  chloride 
forms  with  PtCl4  a  very  difficultly  soluble  complex  salt,  potassium 
platinic  chloride,  K2PtCl6,  into  which  all  the  potassium  in  a  given 
sample  could  be  converted  and  this  form  isolated  and  weighed. 
All  platinum  compounds  are  easily  reduced  to  the  free  ele- 
ment. The  chloride,  for  example,  gives  off  chlorine  upon 
being  heated  to  a  red  heat, 

PtCl4^Pt+2Cl2. 

819.  Gold. — Gold  (Au=i97.2)  has  been  known  from  pre- 
historic times.  It  usually  occurs  in  free  form  (native  gold)  but 
is  also  found  in  compounds  such  as  the  telluride  (815).  Several 
methods  are  in  use  for  the  extraction  of  gold  from  its  ores.  The 
simplest  one  is  the  amalgamation  process,  in  which  sand  or 
crushed  rock  containing  free  gold  is  carried  by  running  water 
over  copper  plates  covered  with  a  film  of  mercury.  The  gold 
forms  an  alloy  (amalgam)  with  the  mercury,  from  which  it  can 
be  freed  subsequently  by  distilling  off  the  latter.  In  the  chlori- 
nation  process  ores  are  extracted  with  a  solution  of  chlorine,  in 
which  gold  dissolves  as  auric  chloride,  AuCl3.  From  the 
chloride  solution  the  gold  may  be  precipitated  by  a  variety 
of  reducing  agents,  e.g.,  ferrous  sulfate,  etc.  More  important 
than  the  preceding  is  the  cyanide  process,  by  which  enormous 
quantities  of  gold  are  recovered  from  low-grade  ores.  This 
process  is  based  on  the  formation  of  readily  soluble  sodium 
auricyanide,  NaAu(NC)2,  which  contains  the  very  stable  com- 
plex ion  Au(NC)2~.  The  reaction  is  as  follows: 

4Au+8NaNC+2H20+03^4NaAu(NC)2+4NaOH. 

The  crushed  ore,  which  must  sometimes  be  roasted  to  free  it 
from  sulfide,  etc.,  is  placed  in  immense  tanks  and  extracted 
with  a  solution  of  sodium  cyanide  (665).  The  resulting  extract 
is  then  run  over  zinc  shavings  to  precipitate  the  gold, 

2NaAu(NC)2-f  Zn->Na2Zn(NC)4+2Au , 


Some  Additional  Elements  and  Their  Compounds         537 

Many  copper,  silver,  and  lead  ores  contain  gold  which 
accompanies  the  metals  when  they  are  obtained  by  smelting 
processes.  This  gold  is  recovered  when  the  metals  are  refined 
electrolytically. 

The  principal  gold-producing  regions  in  the  order  of  their 
importance  are  the  Transvaal  (South  Africa),  United  States 
including  Alaska,  Australia,  Russia  including  Siberia,  Mexico, 
and  Canada.  The  value  of  the  world's  production  of  gold 
during  the  year  1915  was  473  million  dollars. 

Gold  is  a  rather  soft  metal  and  is  usually  alloyed  with  copper 
or  other  metals.  Pure  gold  is  designated  as  24  carat.  Eigh teen- 
carat  gold,  used  in  high-grade  jewelry,  is  75  per  cent  pure. 
United  States  gold  coins  consist  of  90  per  cent  gold  and  10  per 
cent  copper.  Gold  is  the  most  malleable  and  ductile  of  all 
metals. 

Metallic  gold  is  not  attacked  by  any  of  the  common  acids. 
It  dissolves  in  aqua  regia  (562),  a  mixture  of  hydrochloric  and 
nitric  acids,  to  form  auric  chloride,  AuCl3.  This  salt  gives  with 
hydrochloric  acid,  chlorauric  acid,  HAuCl4'3H2O,  forming 
yellow  crystals,  and  with  sodium  chloride,  sodium  chloraurate, 
NaAuCl4'2H2O.  Aurous  chloride,  AuCl,  is  an  insoluble  salt 
obtained  by  gently  heating  AuCl3. 

A  solution  of  auric  chloride  is  used  in  toning  photographs. 
In  the  ensuing  reaction  the  silver  of  the  picture  is  replaced  (492) 
by  metallic  gold  Au, 

AuCl3+3Ag->Au+3AgCl. 

Two  oxides  of  gold  are  known,  aurous  oxide,  Au20,  and  auric 
oxide,  Au2O3;  these  both  give  oxygen  and  metallic  gold  when 
heated. 

820.  The  Rare  Earth  Elements. — The  term  rare  earth  ele- 
ments is  used  to  designate  a  large  group  of  similar  trivalent 
elements.  The  name  indicates  that  these  more  or  less  rare 
elements  bear  some  resemblance  to  aluminum,  the  silicate  of 
which  constitutes  clay,  a  typical  earthy  substance.  One  of 
the  most  abundant  sources  of  these  elements  is  the  mineral 
monazite,  which  is  found  in  North  Carolina,  South  Carolina, 


538  Introduction  to  General  Chemistry 

Florida,  Idaho,  Brazil,  and  India.  Monazite  ordinarily  occurs 
as  a  heavy  yellow  sand.  It  is  a  mixed  phosphate  of  cerium, 
lanthanum,  neodymium,  and  thorium,  with  smaller  proportions 
of  many  other  rare  earth  elements. 

Monazite  sand  is  worked  up  primarily  for  the  thorium  it 
contains.  Strictly  speaking,  thorium  is  not  a  rare  earth  element, 
although  in  its  chemical  behavior  it  is  similar  to  these  elements. 

Thorium  nitrate,  Th(N03)4,  is  extensively  used  to  make 
incandescent  gas  mantles;  these  consist,  to  the  extent  of  over 
99  per  cent,  of  thorium  oxide,  ThO2,  and  must  contain  0.6  to 
0.8  per  cent  of  cerium  oxide,  CeO2  (301). 

A  brief  description  of  the  chemistry  of  cerium,  lanthanum, 
and  neodymium  will  serve  to  illustrate  the  nature  of  the  rare 
earth  elements,  which  are  all  much  alike  in  their  properties. 
In  free  form  the  three  elements  mentioned  are  metals  stable  in 
air  and  toward  water.  An  alloy  of  the  three  metals,  called 
mixed  metal,  is  made  technically  by  the  electrolysis  of  the  molten 
chlorides  CeCl3,  LaCl3,  NdCl3.  An  alloy  of  70  per  cent  mixed 
metal  and  30  per  cent  iron  gives  out  a  shower  of  sparks  of  burning 
metal  when  scratched  with  a  file.  It  is  for  this  reason  known  as 
pyrophoric  alloy  and  finds  use  in  cigar  lighters,  gas  lighters,  etc. 

These  elements  form  bases  and  salts  in  which  they  are 
trivalent.  The  hydroxides,  e.g.,  La(OH)3,  are  almost  insoluble, 
moderately  strong  bases.  The  nitrates,  chlorides,  and  sulfates 
are  soluble  and  are  not  appreciably  hydrolyzed  in  solution.  The 
carbonates,  phosphates,  and  oxalates  are  insoluble  in  water. 
Cerium  differs  from  the  other  two  elements  in  forming  a  series 
of  quadrivalent  salts  derived  from  the  oxide  Ce02.  In  these  the 
eerie  ion,  Ce4+,  is  a  strong  oxidizing  agent.  The  eerie  ion, 
Ce4+,  is  orange  in  color,  while  cerous  ion,  Ce+t+,  is  colorless, 
as  is  also  lanthanum  ion  La+++.  Neodymium  ion,  Nd+"l"f, 
is  a  beautiful  rose  color.  In  addition  to  these  rare  earths 
there  are  eleven  others,  all  of  which  are  less  common  than  the 
three  here  mentioned. 


CHAPTER  XXXI 
CLASSIFICATION  OF  THE   ELEMENTS.    THE   PERIODIC  SYSTEM 

821.  Introduction. — By  reason  of  the  enormous  number  of 
known  chemical  substances  the  best  that  a  trained  chemist  can 
hope  or  even  wish  to  do  is  to  become  thoroughly  familiar  with 
the  principles  and  laws  of  his  science  and  to  know  a  moderate 
number  of  facts  regarding  the  commoner  substances,  together 
with  those  with  which  his  own  field  of  work  brings  him  in  con- 
tact.    His  mind  should  be  a  laboratory  and  not  a  warehouse; 
otherwise  he  will  soon  find  it  so  crowded  with  useless  and  unre- 
lated data  (chemical  junk!)  that  he  has  no  ability  to  solve  the 
new  chemical  problems  that  will  constantly  confront  him.    If  one 
knows  well  the  fundamental  chemical  principles  and  is  familiar 
with  the  chemical  behavior  of  the  most  typical  of  the  elements 
he  will  find  it  easy  to  understand  the  chemistry  of  any  unfamiliar 
element  he  may  have  occasion  to  study.     Moreover,  by  reason 
of  the  relationships  between  the  elements  to  be  discussed  in  this 
chapter  he  will  be  able  from  a  knowledge  of  a  very  few  facts 
regarding  the  element  in  question  to  predict  with  more  or  less 
assurance  the  general  behavior  of  the  element  and  its  com- 
pounds. 

We  shall  now  consider  briefly  the  entire  list  of  the  elements 
with  the  view  of  bringing  out  their  relations  to  one  another 
and  of  showing  that  they  may  be  classified  into  groups  (families) 
and  series  in  such  a  way  as  clearly  to  exhibit  these  relationships. 
In  fact,  every  element  finds  a  definite  place  in  this  system  of 
classification,  which  for  reasons  soon  to  be  shown  is  called 
the  Periodic  System. 

822.  Chemical  Groups  or  Families.    The  Halogens. — The 
similarities  in  properties  and  behavior  of  chlorine,  bromine,  and 
iodine  and  of  their  corresponding  compounds  must  be  well 
known  to  the  reader  at  this  stage  of  advancement.     Let  us 
review  some  of  the  facts  already  established  and  consider  also 

539 


540 


Introduction  to  General  Chemistry 


a  few  additional  ones.  These  elements  are  all  non-metals  which 
form  the  colorless  gases  HC1,  HBr,  and  HI,  all  of  which  dissolve 
easily  in  water  to  form  strong  acids.  In  these  acids  and  their 
salts  the  halogen  is  always  univalent.  Although  the  free  halo- 
gens are  colored  their  ions  in  solution  are  all  colorless.  The  salts, 
with  the  exception  of  those  of  silver,  lead,  and  univalent  mercury, 
are  easily  soluble  in  water.  The  three  halogens  being  con- 
sidered all  form  oxygen  acids  and  salts  of  a  class  represented  by 
KC103.  Thus  we  have  the  acids  chloric,  HC1O3,  bromic, 
HBr03,  and  iodic,  HI03,  and  their  corresponding  salts,  the 
chlorates,  bromates,  and  iodates. 

If  now  we  turn  to  the  differences  between  the  corresponding 
compounds  of  the  three  halogens  we  observe  that  the  properties 
of  bromine  and  its  compounds  are  in  most  cases  nearly  inter- 
mediate between  those  of  chlorine  and  iodine.  Let  us  consider 
first  the  matter  of  atomic  weights : 

Chlorine 35.5 

Iodine 127. 


Mean 81.2 

Bromine 80. 

The  atomic  weight  of  bromine  is  seen  to  be  very  close  to  the 
mean  of  the  atomic  weights  of  the  other  two  elements.  The 
data  shown  in  Table  XLV  illustrate  the  fact  that  bromine  and 

TABLE  XLV 

COMPARISON  OF  THE  PROPERTIES  OF  CHLORINE,  BROMINE,  AND 
IODINE  AND  THEIR  COMPOUNDS 


Chlorine 

Bromine 

Iodine 

Physical  state  
Color  of  gas  or  vapor 

Gas 
Yellow 

Liquid 
Brown 

Solid 
Violet 

Melting-point. 

—  102° 

-     7° 

+  114° 

Boiling-point. 

+  59° 

+  184° 

Density  
Solubility  of  sodium  salt*  
Solubility  of  potassium  salt*  
Melting-point  of  sodium  salt  
Melting-point  of  potassium  salt.  .  .  . 

33 
790° 

3-2 

75 
67 
710° 
750° 

5-o 
180 
140 
650° 
639° 

!  Grams  of  salt  dissolved  by  100  g.  of  water  at  20°. 


Classification  of  the  Elements.     The  Periodic  System      541 


its  compounds  stand  nearly  midway  between  chlorine  and 
iodine  and  their  corresponding  compounds;  so  that  if  we  know 
the  properties  of  compounds  of  the  latter  we  can  predict  pretty 
nearly  those  of  the  former. 

Fluorine,  although  classed  as  a  halogen,  is  less  closely 
related  to  the  three  elements  just  discussed  than  are  the  latter 
to  one  another. 

823.  The  Alkali  Metals. — Sodium  and  potassium  are  called 
the  alkali  metals.  This  group  or  family  also  includes  lithium 
(786),  a  much  less  common  element,  and  in  addition  two  rare 
elements,  rubidium  and  caesium.  These  two  elements  form 
with  sodium  and  potassium  as  closely  related  a  group  as  do 
chlorine,  bromine,  and  iodine.  Lithium,  although  showing 
many  points  of  similarity  to  the  other  four  elements,  bears  about 
the  same  sort  of  relation  to  the  others  that  fluorine  does  to  the 
other  halogens.  Table  XL VI  gives  some  typical  data  on  this 

TABLE  XLVI 

COMPARISON  OF  THE  PROPERTIES  OF  THE  ALKALI  METALS  AND  THEIR  COMPOUNDS 


Lithium 

Sodium 

Potassium 

Rubidium 

Caesium 

Atomic  weights  
Melting-points  

6.94 

186° 

23.00 
98° 

39.01 
62° 

85.45 
38° 

132.81 
26° 

Densities  

0.59 

0.97 

0.87 

1-52 

1.88 

A.tomic  volumes 

119 

23  .  7 

44.8 

56.1 

70.6 

Hydroxides 

LiOH 

NaOH 

KOH 

RbOH 

CsOH 

Chlorides  
Nitrates  

LiCl 
LiN03 

NaCl 
NaNO3 

KC1 
KN03 

RbCl 
RbNO3 

CsCl 
CsNO3 

group.  Again  we  find  some  simple  relations  between  atomic 
weights,  as  in  the  case  of  the  halogens.  Thus  the  atomic  weight 
of  sodium,  23,  is  practically  the  mean  of  the  atomic  weights  of 
lithium  and  potassium,  while  that  of  rubidium  stands  almost 
midway  between  those  of  potassium  and  caesium.  Curiously 
enough,  however,  the  atomic  weight  of  potassium,  39,  is  much 
less  than  the  mean  of  the  values  for  the  neighboring  elements 
sodium  and  rubidium,  namely  54. 

All  of  the  five  elements  of  this  family  are  typical  metals. 
As  such  they  are  all  very  active,  caesium  most  so  and  lithium 
least.  With  water  they  all  react  readily  to  form  strong,  soluble 


542  Introduction  to  General  Chemistry 

bases,  resembling  sodium  hydroxide.  This  means  that  they  all 
form  positive  ions,  all  of  which  are  univalent.  The  metals  are 
all  soft  (easily  cut  with  a  knife).  They  are  all  silver- white  and 
have  low  melting-points,  rubidium  and  caesium  melting  below 
blood  heat.  The  melting-points  decrease  in  order  from  lithium 
to  caesium.  The  densities  are  all  lower  than  those  of  any  other 
metals.  In  general  the  higher  the  atomic  weight  the  greater 
the  density;  but  we  find  that  the  value  for  sodium  is  exceptional 
in  being  a  little  higher  than  that  of  potassium.  If,  however,  we 
calculate  from  the  density  and  atomic  weight  the  so-called  atomic 
volume,  that  is  the  volume  in  cubic  centimeters  occupied  by  one 
gram  atomic  weight  of  the  element,  we  find  that  these  constants 
increase  steadily  with  increasing  atomic  weight,  and  that  the 
value  for  sodium  is  no  longer  irregular.  Since  the  gram  atomic 
weights  (symbol  weights)  of  all  elements  contain  equal  numbers 
of  atoms,  the  atomic  volumes  are  the  volumes  of  equal  numbers 
of  atoms.  By  reference  to  Table  XL VI  we  see  that  in  this  family 
the  heavier  an  atom  is  the  greater  the  volume  it  occupies. 

824.  The    Alkali    Earth    Group.— The    alkali    earth    group 
includes  the  four  elements  calcium,   strontium,   barium,   and 
radium.     These  are  all  bivalent  metals  whose  hydroxides  are 
moderately  strong  bases,  and  whose  salts  are  typically  repre- 
sented by   calcium   chloride,    CaCl2,   and   sulfate,  CaSO4.    A 
detailed  examination  of  this  group  would  show  that  its  members 
closely  resemble  one  another  in  general,  and  also  that  a  system- 
atic and  gradual  change  of  properties  accompanies  change  of 
atomic  weight  just  as  in  the  halogen  and  alkali  metal  families. 

825.  The  Inert  Gases. — The  inert  gases  helium,  neon,  argon, 
krypton,  and  xenon  (791-797)  must  obviously  be  considered  as 
belonging  to  a  separate  family.    They  are  all  devoid  of  chemical 
properties  and  are  therefore  of  valence  zero  in  all  cases.    As 
gases  their  densities  are  of  course  all  proportional  to  their 
molecular  weights  (Avogadro's  Law,  193).    The  ratio  of  the 
molecular  heat  at  constant  pressure  to  that  at  constant  volume 
is  i .  66  in  each  case,  thus  showing  that  each  gas  is  monatomic 
(793).    The   critical   temperatures  and  boiling-points   of   the 
liquid  gases  all  increase  with  increasing  atomic  weight  from  the 


Classification  of  the  Elements.     The  Periodic  System     543 


lowest  values  found  in  the  case  of  helium.  No  other  elements 
(excepting  niton,  radium  emanation,  chap,  xxxii)  resemble  the 
five  members  of  this  family  at  all  closely.  Sometimes  the  inert 
gases  are  called  the  noble  gases,  because  the  most  striking  char- 
acteristic of  the  noble  metals,  gold  and  platinum,  is  the  reluctance 
with  which  they  enter  into  chemical  combination. 

826.  Series  of  Elements. — If  we  arrange  the  elements  in  the 
order  of  their  atomic  weights  the  first  of  the  series  is  obviously 
hydrogen.  The  next  is  helium.  The  next  seven  are  shown  in 
Table  XL VII.  The  data  presented  in  this  table  show  that  the 

TABLE  XL VII 

THE  FIRST  SERIES  OF  ELEMENTS 


Li 

Be 

B 

C 

N 

0 

F 

Serial  numbers  .... 

? 

4' 

c 

6 

7 

8 

Atomic  weights  
Melting-points  

7 
1  86° 

1800° 

II 
2  3  SO0 

12 

3600° 

14 
-210° 

16 
-218° 

19 
-223° 

Chlorine  compounds  

LiCl 

BeCl2 

BCL 

CCL 

Valence  toward  chlorine  .  . 
Hydrogen  compounds  .... 
Valence  toward  hydrogen  . 

i 

2 

3 

4 
CH4 

4 

NH3  ' 

T. 

H20 

2 

HF 

i 

valence  toward  chlorine  increases  by  one  as  we  pass  from  lithium, 
one,  to  carbon,  four,  and  that  the  valence  toward  hydrogen 
decreases  by  units  from  carbon,  four,  to  fluorine,  one.  There  is 
also  a  gradual  change  of  both  physical  and  chemical  properties 
as  we  go  from  lithium  to  fluorine.  For  example,  lithium  is  a 
light,  soft,  silver- white  metal  (density  o .  59) .  It  readily  decom- 
poses cold  water.  Lithium  hydroxide  is  a  strong  base,  and  its 
salts  are  not  hydrolyzed  in  solution.  Beryllium  is  a  white 
malleable  metal  of  density  i .  64.  It  does  not  tarnish  in  air, 
but  it  slowly  decomposes  hot  water.  Beryllium  hydroxide  is 
a  rather  weak  base.  The  chloride  is  appreciably  hydrolyzed  in 
water  solution.  Boron  (801-803)  is  a  non-metallic  solid  of 
very  high  melting-point.  It  is  not  acted  on  by  water  or  by 
oxygen  except  at  high  temperatures.  Its  hydroxide,  B(OH)3, 
or  H3B03,  boric  acid,  is  an  extremely  weak  acid.  The  chloride 
BC13  is  completely  hydrolyzed  by  water  to  form  boric  and 


544 


Introduction  to  General  Chemistry 


hydrochloric  acids.  Carbon  is  strictly  non-metallic  in  the  form 
of  diamond  (the  latter  is  an  electrical,  insulator-like  glass),  and 
only  semimetallic  as  graphite  (this  has  a  semimetallic  luster  and 
conducts  the  current  fairly  well,  630).  Carbon  tetrachloride, 
CC14  (644),  is  a  colorless  liquid  with  none  of  the  characteristics 
of  a  salt.  The  other  three  members  of  the  series,  nitrogen, 
oxygen,  and  fluorine,  are  all  gases  and  all  strictly  non-metallic 
in  character.  Their  activity  increases  very  markedly  from 
nitrogen,  through  oxygen,  to  fluorine.  In  their  hydrogen  com- 
pounds the  last  four  elements  of  the  series  show  a  progressive 
decrease  in  valence  from  four  for  carbon  to  one  for  fluorine. 
Only  the  hydrogen  compounds  of  nitrogen,  oxygen,  and  fluorine 
form  compounds  with  metals,  e.g.,  NaNH2  (527),  NaOH,  and 
NaF.  The  first  is  decomposed  completely  by  water. 

827.  The  Second  Series.— The  seven  elements  beginning 
with  sodium  taken  in  increasing  order  of  their  atomic  weights 
form  a  second  series  as  illustrated  in  Table  XL VIII.  Again  in 

TABLE  XL VIII 

THE  SECOND  SERIES  OF  ELEMENTS 


j      Na 

Mg 

Al 

Si 

P 

S 

Cl 

Serial  numbers  1  1 

12 

17 

14 

1C 

16 

17 

Atomic  weights  23 

24.3 

27 

28 

3i 

32 

35  -5 

Chlorine  compounds            !  NaCl 

MgCla 

A1C13 

SiCl4 

PCL 

Valence  toward  chlorine  .  .        i 

2 

3 

4 

5 

Hydrogen  compounds  
Valence  toward  hydrogen  

SiH4 

4 

PH3 

3 

H2S 

2 

HC1 

i 

this  series  metallic  properties  characterize  the  first  three  elements. 
The  fourth,  silicon,  is  a  hard,  black,  difficultly  fusible  element 
with  fair  electrical  conductivity;  it  may  well  be  classed  as  semi- 
metallic.  The  following  three  elements  are  truly  non-metallic. 
Chemical  activity,  great  in  the  case  of  sodium,  diminishes  toward 
silicon,  which  is  quite  inert,  and  then  increases  again  to  a  maxi- 
mum with  chlorine.  In  this  connection  it  is  of  importance  to 
note  that  the  activities  of  the  first  three  metals  are  represented 
by  their  tendencies  to  take  on  positive  charges  (lose  electrons, 
491),  while  those  .of  the  last  three  elements  are  due  to  a  tendency 


Classification  of  the  Elements.    The  Periodic  System     545 


to  acquire  negative  charges  (gain  electrons,  489).  Furthermore 
the  valences  change  by  unity  as  we  pass  from  one  to  the  next 
element  in  the  series. 

828.  Correlation  of  the  First  and  Second  Series. — If  we 
write  the  two  series  of  elements  in  parallel  columns  a  very 
remarkable  relationship  at  once  appears,  as  shown  in  Table 
XLIX.  We  see  that  Li  and  Na,  both  members  of  the  alkali- 

TABLE  XLIX 


First  series    

Li 

Be 

B 

C 

N 

o 

F 

Second  series  

Na 

Mg 

Al 

Si 

P 

S 

Cl 

H  compound  

EH4 

EH3 

EH2 

EH 

Cl  compound 

EC1 

ECL 

ECL 

ECL 

ECL 

metal  group,  fall  in  the  same  vertical  column;  likewise  the 
halogens,  F  and  Cl.  Furthermore,  if  E  is  written  as  the  general 
symbol  of  any  element,  the  compounds  with  H  or  Cl  for  the  two 
elements  of  any  vertical  column  can  be  represented  by  the  same 
formula.  In  other  words,  this  arrangement  brings  elements 
of  similar  properties  into  the  same  vertical  columns. 

829.  The  Periodic  Table  of  the  Elements. — Table  L  shows 
a  systematic  arrangement  of  70  of  the  83  elements.  In  this 
table  the  first  two  horizontal  lines  reproduce  the  two  series  of  the 
preceding  section  and  in  addition  include  the  elements  helium,  He, 
and  neon,  Ne.  Following  chlorine  all  known  elements  to  and 
including  cerium,  Ce,  140.2,  are  given  in  the  order  of  increasing 
atomic  weights,  with  the  exception  that  argon,  A,  39 . 9,  precedes 
potassium,  K,  39 .  i ;  and  tellurium,  Te,  127.5,  precedes  iodine,  I, 
126.9.  Between  cerium,  Ce,  140.2,  and  tantalum,  Ta,  181 .5,  a 
number  of  rare  earth  elements  are  omitted  (820,  846). 

The  first  and  second  series,  called  also  the  first  and  second 
periods  respectively,  each  includes  eight  elements.  The 
double  series  of  18  elements  beginning  with  argon,  A,  39.9, 
and  ending  with  bromine,  79.9,  forms  the  third  period.  It 
will  be  noticed  that  a  group  of  three  elements,  iron,  Fe,  cobalt, 
Co,  and  nickel,  Ni,  follows  manganese,  Mn,  54.9,  in  the  column 
headed  Group  VIII.  Copper,  the  next  element  after  nickel,  is 
placed  in  Group  I  and  not  under  argon  in  Group  O.  The 


546 


Introduction  to  General  Chemistry 


a3 


a 


3a 


•-•a 
ex 


O\ 


Group  IV 
RO,  RH 


H  H 


o 


:(S 


85 


saoiaad  XHOHS 


saoiaad  ONOI 


Classification  of  the  Elements.     The  Periodic  System      547 

fourth  period  begins  with  krypton,  Kr,  82.9,  and  ends  with 
iodine,  126 .9.  The  fifth  period  includes  the  elements  beginning 
with  xenon,  X,  130.2.  The  sixth  period,  which  is  defective, 
shows  tantalum  as  its  first  and  bismuth  as  its  last  member. 
The  seventh  period  contains  only  four  elements,  of  which  niton, 
Nt,  222.4,  is  the  first  and  uranium,  U,  238.2,  the  last.  No 
element  of  higher  atomic  weight  than  uranium  is  known.  The 
first  two  periods  are  called  short  periods,  and  the  next  four  long 
periods.  The  seventh,  and  last,  period  is  fragmentary. 

830.  The  Nine  Groups  of  Elements.    The  Zero  Group.— The 
elements  contained  in  any  one  of  the  nine  vertical  columns 
headed  O  to  VIII  constitute  a  group.     There  are,  therefore,  nine 
such  groups.     The  zero  group  consists  of   the  five  inert  gases 
(rare  or  noble  gases  studied  earlier,  825),  together  with  niton 
(radium  emanation,  see  chap,  xxxii),  which  is  also  inert  chemi- 
cally.   None  of  these  gases  forms  any  chemical  compounds.    The 
members  of  the  zero  group  never  have  other  than  zero  valence. 

831.  The  First  Group.    A  and  B  Families. — Reference  to  the 
Periodic  Table  shows  that  five  of  the  eight  elements  of  Group  I 
belong  to  the  alkali-metals  family.    These  elements  are  lithium, 
sodium,  potassium,  rubidium,  and  caesium.    The  other  three 
elements,  copper,  silver,  and  gold,  as  we  already  know  from 
their  earlier  study,  are  widely  different  from  sodium  and  potas- 
sium, the  best-known  alkali  metals.     In  many  respects  copper, 
silver,  and  gold  are  much  alike  and  may  be  classed  together  as 
a  family.     Let  us  compare  and  contrast  the  two  families  of 
Group  I,  which  we  may  designate  as  I A  and  IB  respectively. 
The  differences  are  set  forth  in  Table  LI. 

TABLE  LI 


ELEMENTS 

IA 

IB 

Li,Na,K,Rb,Cs 

Cu,  Ag,  Au 

Density 

o  .  59  to  i  .  88 
Very  soft 
26°  to  1  86° 
Give  hydroxides 
i  Soluble  strong  bases 
Very  high 

8  .  8  to  19  .  2 
Hard 
960°  to  1060° 
No  action 
Insoluble  or  unstable 
Very  low 

Hardness 

Melting-point                          .        .  . 

Action  of  Water  
Hydroxide 

Displacement  power  (402) 

548  Introduction  to  General  Chemistry 

The  differences  are  set  forth  in  Table  LI.  The  members  of  IA 
are  uniformly  univalent  when  in  combination  and  form  such 
compounds  as  EC1,  EBr,  EN03,  E2S04,  EOH,  and  E20.  In 
sodium  peroxide,  Na2O2,  we  believe  that  sodium  is  still  univalent 
and  represents  the  structural  formula  thus:  Na  •  0  •  O  •  Na  (324) . 
In  IB  copper  and  gold  are  variable  in  their  valences;  thus  copper 
forms  in  addition  to  the  commoner  compounds  in  which  it  is 
bivalent  (165)  a  series  of  univalent  compounds,  of  which  cuprous 
chloride,  CuCl,  is  an  example.  Gold  has  so  little  chemical 
affinity  that  it  does  not  form  many  stable  compounds,  but  in 
these  it  is  usually  univalent  or  trivalent.  It  forms  two  oxides, 
Au2O  and  Au2O3.  If  we  fix  our  attention  on  the  resemblances 
between  copper,  silver,  and  gold  rather  than  on  the  divergences 
it  is  easy  to  see  why  they  should  be  classed  in  the  same  family. 
The  metals  themselves  are  all  permanent  in  air  and  not  very 
active  chemically;  none  of  them  sets  free  hydrogen  from  acids 
appreciably.  They  are  moderately  hard  but  are  very  malleable 
and  ductile.  They  all  melt  in  the  neighborhood  of  1000°.  They 
are  the  best  three  conductors  of  heat  and  of  electricity.  They 
all  form  oxides,  E2O,  and  their  chlorides,  EC1,  are  all  white  salts 
insoluble  in  water  and  dilute  acids.  It  will  be  noted  that  Cu, 
Ag,  and  Au  occur  in  the  long  periods  3,  4,  and  6  respectively, 
and  in  each  case  in  the  second  line  of  the  period.  Inspection 
will  show  the  presence  of  A  and  B  families  in  Groups  II  to  VII 
but  not  in  groups  O  and  VIII. 

832.-  The  Second  Group. — Group  II  is  made  up  of  two 
families,  IIA  and  IIB.  The  alkali  earth  elements  Ca,  Sr,  Ba, 
and  Ra  are  the  typical  members  of  IIA  (824);  while  Zn,  Cd 
(cadmium),  and  Hg  constitute  IIB.  Beryllium,  Be, .and  mag- 
nesium, Mg,  in  the  first  and  second  periods  are  somewhat 
more  closely  related  in  their  properties  to  the  family  IIA 
than  to  the  family  IIB.  The  alkali  earth  family,  IIA,  is 
so  called  because  of  the  resemblance  of  its  members  to 
the  alkalies,  IA,  on  the  one  hand,  and  to  the  earths,  typified 
by  aluminum  in  Group  III,  on  the  other.  Metallic  .calcium 
is  a  moderately  soft  metal  and  is  acted  on  rather  rapidly  by 
water  to  form  a  strongly  basic  hydroxide  and  hydrogen,  but 


Classification  of  the  Elements.     The  Periodic  System      549 

by  no  means  so  rapidly  as  are  the  alkali  metals.  The  hydroxides 
of  IIA  are  not  quite  as  strong  bases  as  those  of  IA;  but  their 
salts  with  strong  acids  give  neutral  solutions.  The  three  mem- 
bers of  IEB,  Zn,  Cd,  and  Hg,  are  typical  metals  and  are  all  unacted 
upon  by  water  and  are  all  untarnished  by  air.  The  hydroxides 
of  IIB  are  weak,  insoluble  bases,  and  solutions  of  salts  with  strong 
acids  are  acid  in  reaction  by  reason  of  hydrolysis  (436).  Com- 
parison of  IIA  and  IIB  shows  less  difference  in  properties  than 
was  found  in  the  case  of  I A  and  IB.  With  the  exception  of 
mercury  all  the  elements. of  IIA  and  IIB  form  exclusively  com- 
pounds in  which  they  are  bivalent.  Typical  formulae  are  EC12, 
EBr2,  E(N03)2,  ES04,  EC03,  EO,  and  E(OH)2.  Mercury  is 
exceptional,  since  in  mercurous  compounds  (333)  it  is  univalent. 

833.  The  Third  Group. — Of  the  nine  elements  composing  the 
third  group,  only  two,  boron  (801)  and  aluminum  (174),  are 
common.    The  others,  with  the  exception  of  lanthanum,  are 
rare.    In  their  compounds  the  third-group  elements  are  always 
trivalent  and  give  products  of  the  following  types :  E2O3,  E(OH)3, 
EC13,  E(N03)3,  E2(S04)3. 

Boron  is  the  only  member  of  the  group  that  does  not  form 
salts  with  acids;  its  hydroxide,  boric  acid  (802),  has  only  weak 
acid  properties.  Aluminum  hydroxide,  A1(OH)3,  is  amphoteric 
(177),  forming  salts  with  both  acids  and  bases;  but  these  salts 
are  all  considerably  hydrolyzed  in  solution,  so  that,  for  example, 
a  solution  of  A1C13  reacts  acid,  while  one  of  NaAlO2  reacts  basic, 
to  indicators  (436). 

Lanthanum  hydroxide,  La(OH)3,  a  white  solid,  is  a  rather 
strong  base  forming  salts,  e.g.,  La(NO3)3,  which  are  not  hydro- 
lyzed in  solution. 

The  rare  earths  (820),  which  doubtless  should  be  considered 
as  third-group  elements,  will  be  treated  separately  in  a  later 
section  (846). 

834.  The  Fourth  Group. — All  the  elements  of  the  fourth 
group,  with  the  exception  of  germanium  (Ge  =  72 .5),  are  of  use 
technically.     Carbon  and  silicon  are  of  course  the  most  impor- 
tant, and  tin  and  lead  come  next.    Titanum,  zirconium,  cerium, 
and  thorium  find  interesting  minor  uses.    The  characteristic 


550  Introduction  to  General  Chemistry 

valence  of  these  elements,  in  their  compounds,  is  four,  so  that 
the  typical  compounds  have  formulae  like  the  following:  E02, 
E(OH)4,  H2E03,  EC14,  E(NO3)4,  E(S04)2,  etc.  Only  titanium, 
zirconium,  tin,  cerium,  and  thorium  form  salts  of  the  last  two 
types.  Several  of  the  members  of  this  group  form  compounds 
in  which  they  have  a  valence  of  two  or  three.  Thus  carbon  and 
lead  form  the  oxides  CO  (632)  and  PbO  (167)  respectively.  In 
all  stannous  compounds  (809)  tin  is  bivalent,  while  in  all  its 
ordinary  salts  lead  also  has  a  valence  of  two.  In  the  more  stable 
salts  of  cerium,  like  cerous  nitrate,  Ce(N03)3,  and  cerous  sulfate, 
Ce2(S04)3,  this  element  is  trivalent;  but  moist  cerous  hydroxide, 
Ce(OH)3,  takes  up  oxygen  from  the  air  to  form  the  more  stable 
eerie  hydroxide,  Ce(OH)4.  All  the  fourth-group  elements  form 
chlorides  of  the  type  EC14.  All  of  these  except  CeCl4  and  PbCl4 
are  stable  toward  heat;  the  chlorides  of  carbon,  silicon,  titanium, 
germanium,  and  tin  are  colorless  liquids  of  low  boiling-points 
(59°  to  135°). 

The  hydroxides  of  the  elements  of  the  fourth  group  form  only 
weak  acids  or  weak  bases.  Carbon  and  silicon  give  carbonic  acid 
and  silicic  acid  (807)  respectively;  and  titanium,  zirconium, 
cerium,  and  thorium  form  hydroxides  which  are  weak  bases. 
Stannous  and  stannic  hydroxides  (809)  and  lead  hydroxide, 
Pb(OH)2,  are  amphoteric.  Titanium,  zirconium,  and  thorium 
are  the  typical  members  of  IVA;  the  position  of  cerium  in  this 
group  is  a  bit  uncertain.  Germanium,  tin,  and  lead  compose 
IVB.  Carbon  and  silicon  bear  resemblances  to  both  of  these 
families.  Tin  and  lead  are  easily  obtained  as  free  metals,  but 
it  is  doubtful  whether  even  moderately  pure  metallic  thorium 
has  ever  been  obtained. 

835.  The  Fifth  Group.— We  have  studied  only  five  of  the 
eight  elements  given  in  Group  V,  Table  L.  These  are  nitrogen, 
chapter  xxi;  phosphorus,  chapter  xxiii;  arsenic  (810),  antimony 
(811),  and  bismuth  (812).  The  last  three  form  family  VB,  to 
which  the  first  two  show  points  of  similarity.  Of  the  members 
of  family  VA  only  vanadium  (V=5i)  is  of  technical  importance. 
It  is  a  rather  scarce  element  which  finds  an  important  use  as  a 
minor  component  of  high-grade  machine  steel.  Its  common 


Classification  of  the  Elements.     The  Periodic  System      551 

oxide  is  V2OS,  which  is  weakly  basic  and  rather  strongly  acidic 
in  character.  The  metavanadates,  e.g.,  NH4V03,  correspond 
to  the  salts  of  metaphosphoric  acid,  HPO3  (589).  No  technical 
uses  have  yet  been  found  for  the  other  two  members  of  this 
family,  the  rare  elements  columbium  (also  called  niobium)  and 
tantalum. 

The  following  discussion  will  be  confined  to  the  members  of 
VB,  together  with  nitrogen  and  phosphorus.  Of  the  five  elements 
referred  to,  all  but  bismuth  form  gaseous  hydrogen  compounds 
of  the  type  EH3.  Of  these  only  NH3  forms  with  water  a  basic 
hydroxide.  In  these  hydrides  the  fifth-group  element  is  tri- 
valent.  All  five  of  the  elements  (including  bismuth)  form 
trimethyl  derivatives,  E(CH3)3,  corresponding  to  trimethyl 
amine  (59). 

All  five  elements  form  trichlorides,  EC13,  and  P,  As,  and  Sb 
form  pentachlorides,  EC1S.  All  of  the  five  elements  being  con- 
sidered form  trioxides,  E2O3,  and  all  with  the  possible  exception 
of  bismuth  form  pentoxides,  E2OS.  Also,  with  the  exception  of 
bismuth,  these  elements  form  acids  derived  from  their  pentox- 
ides. These  acids  exist  in  various  stages  of  hydration,  of  which 
HE03  is  the  most  common. 

The  highest  valence  of  the  elements  of  the  fifth  group  is  five, 
although  in  many  cases  the  valence  is  only  three. 

836.  The  Sixth  Group.— The  two  families  VIA  and  VIB  are 
each  made  up  of  four  elements.  The  first,  VIA,  consists  of 
chromium  (344),  molybdenum  (813),  tungsten  (814),  and 
uranium;  the  second,  VIB,  embraces  oxygen,  sulfur  (chap,  xxiv), 
selenium  (815),  and  tellurium  (815).  The  members  of  VIA  have 
a  characteristic  maximum  valence  of  six,  as  exemplified  by  the 
oxides,  EO3.  These  are  all  acidic  in  nature  and  lead  to  salts 
such  as  Na2EO4  and  Na2E2O7  and  even  more  complex  formulae^ 
e.g.,  Na6Mo7O24  •  22H20.  In  all  these  salts  the  sixth-group 
element  is  hexavalent.  The  members  of  family  VIB  all  form 
hydrogen  compounds,  H2E,  and  these  decrease  in  stability  as  the 
atomic  weight  of  the  sixth-group  element  increases.  Water  is 
extremely  stable:  hydrogen  telluride  decomposes  with  ease  (815). 
We  should  expect  a  normal  valence  of  six  for  these  elements; 


552  Introduction  to  General  Chemistry 

but  we  find  that  for  oxygen  the  ordinary  valence  is  only  two, 
although  in  some  cases  (not  discussed  in  this  book)  a  valence 
of  four  is  very  probable.  Sulfur  and  tellurium  form  trioxides, 
EO3,  and  in  these  the  two  former  are  undoubtedly  hexavalent. 
as  they  are  also  in  the  corresponding  acids  H2E04  and  their  salts, 
While  selenium  trioxide  has  not  been  made,  the  corresponding 
acid  H2Se04  and  its  salts  are  well  known  (815).  In  them 
selenium  has  a  valence  of  six.  Sulfur,  selenium,  and  tellurium 
also  form  oxides  EO2  and  acids  H2EO3  and  their  salts.  These,  as 
we  have  learned,  are  converted  by  oxidation  into  the  correspond- 
ing acids  H2EO4  or  their  salts. 

Contrasting  VIA  with  VIB,  we  may  say  that  while  all  the 
elements  are  acid-forming,  only  chromium  and  uranium  are 
base-forming  and  give  salts  with  acids.  Chromium  forms 
hydroxides  in  which  it  is  trivalent.  Thus  Cr(OH)3  is  a  weak 
base  like  Fe(OH)3.  It  also  is  capable  of  further  reduction  and 
gives  a  more  strongly  basic  hydroxide,  Cr(OH)2.  The  com- 
monest uranium  salts  are  derived  from  the  trioxide  UO3.  These 
salts  may  be  considered  as  being  formed  from  a  compound 
U02(OH)2,  a  diacid  base  about  as  strong  as  Fe(OH)2.  Its 
reaction  with  nitric  acid  may  be  represented  thus: 

UO2(OH)2+2HN03->U02(N03)2+2H2O . 

This  salt  is  called  uranyl  nitrate.  It  forms  large  yellow  crystals, 
U02(N03)2  •  6H2O.  The  corresponding  uranyl  chlorides,  sul- 
fates,  acetates,  etc.,  are  readily  prepared.  In  these  salts 
uranium  is  undoubtedly  hexavalent. 

837.  The  Seventh  Group. — Only  one  member  of  the  A  family 
of  the  seventh  group  is  known;  this  is  manganese  (342).  It  is 
a  hard  metal  the  color  of  iron  and  does  not  tarnish  in  air.  It 
does  not  form  a  hydride.  In  its  common  halogen  compounds 
(true  salts)  it  is  bivalent.  Its  corresponding  hydroxide  is 
Mn(OH)2,  an  insoluble  base  which  forms  salts  like  MnBr2, 
Mn(N03)2,  MnS04,  etc.,  and  an  oxide,  MnO.  Manganese  also 
gives  higher  oxides,  some  of  which  are  acid  anhydrides.  Potas- 
sium manganate,  K2MnO4,  soluble  green  crystals,  is  a  salt  of  the 
(unstable)  manganic  acid  H2MnO4,  the  anhydride  of  which  would 


Classification  of  the  Elements.     The  Periodic  System      553 

be  MnO3.  By  oxidation  (with  chlorine,  for  example)  solutions 
of  manganates  give  permanganates, 

2K2MnO4+Cl2->2KCl+2KMnO4 . 

Potassium  permanganate  (343)  by  treatment  with  sulfuric  acid 
gives  an  explosive  liquid  (danger!)  which  is  very  probably  man- 
ganese heptoxide,  Mn2O7,  the  anhydride  of  permanganic  acid, 
HMnO4.  In  the  latter  and  its  salts  and  anhydride  manganese 
has  a  valence  of  seven.  These  compounds  are  all  violent 
oxidizing  agents. 

The  halogen  family,  VIIB,  is  so  well  known  and  the  close 
resemblances  of  its  members  have  been  so  often  referred  to  that 
further  discussion  of  them  would  be  superfluous.  A  few  lines 
may  be  added  touching  the  valence  of  chlorine,  bromine,  and 
iodine  in  their  oxyacids  and  salts.  Fluorine,  it  will  be  recalled, 
does  not  form  such  compounds.  The  other  three  halogens 
reach  their  maximum  oxygen  valence  in  the  acids  perchloric, 
HC1O4,  perbromic,  HBrO4,  and  periodic,  HI04.  These  may  be 
considered  as  derived  from  hypothetical  heptoxides,  E2O7,  in 
which  the  halogen  has  a  positive  valence  of  seven.  There  is 
thus  an  analogy  between  these  elements  and  manganese  in  the 
peracids  and  their  salts. 

838.  The  Eighth  Group. — An  inspection  of  the  Periodic  Table 
(829)  will  show  that  the  arrangement  of  the  members  of  the 
eighth  group  is  different  from  that  in  any  other  group.  In  this 
group  there  are  three  lines  of  three  elements  each  in  place  of  the 
usual  A  and  B  columns.  In  first  line  we  find  iron,  cobalt,  and 
nickel,  which  are  the  only  common  elements  of  Group  VIII. 
These  elements  have  many  similar  properties.  First  of  all  it 
will  be  noted  that  their  atomic  weights  are  all  close  together. 
These  three  elements  are  all  metals  and  all  form  salts  of  the  types 
EC12,  E(NO3)2,  ESO4,  E3(P04)2,  and  hydroxides  and  oxides 
E(OH)2  and  EO  respectively.  In  all  these  compounds  these 
three  elements  are  bivalent.  However,  iron  and  to  a  lesser 
extent  cobalt  form  hydroxides,  oxides,  and  salts  in  which  they 
are  trivalent,  e.g.,  E(OH)3,  E2O3,  and  EF3.  The  second  line  of 
the  eighth  group  contains  the  three  metals  ruthenium,  Ru; 


554  Introduction  to  General  Chemistry 

rhodium,  Rh,  and  palladium,  Pd,  all  rare  elements  which  closely 
resemble  one  another.  Recently  an  alloy  of  palladium  with  gold 
has  come  into  use  as  a  substitute  for  platinum  in  the  manufacture 
of  crucibles  and  dishes  for  laboratory  use.  This  alloy  closely 
resembles  platinum  in  its  physical  properties  and  in  its  inertness 
toward  chemical  reagents.  The  elements  of  the  third  line 
osmium,  Os,  iridium,  Ir,  and  platinum,  Pt,  form  another  sub- 
group in  which  the  three  members  are  much  alike.  All  are 
extremely  resistant  to  attack  by  most  chemical  reagents. 
Metallic  osmium  is  of  interest  in  that  it  has  the  greatest  density 
of  any  known  substance,  namely  22.5,  It  is  also  extremely 
hard  and  melts  only  at  the  very  high  temperature  of  2500°. 
It  is  also  unique  in  forming  an  easily  volatile  and  extremely 
poisonous  oxide,  Os04,  a  solution  of  which,  known  as  osmic  acid, 
is  a  very  important  staining  material  for  microscopic  prepara- 
tions. In  this  oxide  osmium  has  a  valence  of  eight. 

839.  Valence  and  the  Structure  of  Inorganic  Molecules. — 
The  study  of  organic  chemistry  has  made  it  very  clear  that  the 
valence  of  each  element  of  a  compound  can  be  definitely  deter- 
mined only  when  the  structure  (648)  of  the  molecule  of  the  sub- 
stance is  known.  'It  is  much  easier  to  discover  the  structural 
formulae  of  organic  compounds  than  of  inorganic.  There  are 
two  reasons  for  this:  first,  a  single  element,  carbon,  forms  the 
backbone,  so  to  speak,  of  all  organic  substances;  and  second, 
the  enormous  number  of  carbon  compounds  makes  it  possible 
to  test  theory  by  fact  in  a  multitude  of  cases.  The  problem  is 
very  different  with  inorganic  compounds,  where  we  must  deal 
with  80  or  more  elements,  any  one  of  which  (excepting  hydrogen, 
oxygen,  and  nitrogen)  forms  but  few  compounds  in  comparison 
with  the  host  of  carbon  derivatives.  Therefore  when  we  attempt 
to  show  the  structure  of  inorganic  molecules  the  element  of  un- 
certainty is  often  great. 

In  writing  graphic  formulae  for  the  chlorides,  oxides,  hy- 
droxides, etc.,  of  the  elements  of  the  first  three  groups  we  are 
on  pretty  safe  ground,  as  these  formulae  are  very  simple.  For 
common  salt  we  have  only  one  choice,  Na-  Cl;  similarly  for  other 
univalent  elements.  If  for  water  (323)  we  write  H-OH,  for 


Classification  of  the  Elements.     The  Periodic  System      555 

sodium  hydroxide  we  must  write  Na-O-H.  Sodium  oxide, 
Na2O,  must  be  Na'ONa.  In  the  second  and  third  groups  the 
task  is  nearly  as  easy.  Calcium  chloride  is  of  course  Cl-  Ca-  Cl, 
and  the  oxide  must  be  simply  Ca=O,  calcium  and  oxygen  both 
being  bivalent.  Calcium  hydroxide  then  is  represented  by 
H-OCa-OH. 

In  the  third  group  boron,  which  forms  a  trichloride,  BC13,  is 
evidently  trivalent.  If  boric  acid,  H3BO3,  is  written 

H(X 
HO— B 
HO/ 

the  normal  valences  of  all  three  elements  are  correctly  repre- 
sented. Metaboric  acid,  HBO2,  would  then  be  HO-B  =  O  and 
sodium  metaborate  (802)  NaOB  =  0.  Sodium  aluminate, 
NaAlO2  (177),  would  then  be  written  NaO-Al  =  0,  which  shows 
aluminum  with  its  correct  valence  of  three. 

The  compound  C(OH)4  is  not  known,  but  carbonic  acid, 
which  would  result  from  this  by  loss  of  water,  is  doubtless  cor- 
rectly represented  by 

HOV 

>C=O 

HO/ 

which  shows  carbon  with  a  valence  of  four.  The  structure  of 
silicic  acid  is  probably  analogous. 

Nitrogen  pentoxide,  N2OS  (555),  is  the  anhydride  of  nitric 
acid.  We  might  expect  a  hydroxide  N(OH)S  or  HSNO5.  By 
loss  of  water  this  could  form  first  H3NO4  and  then  HN03.  As 
a  matter  of  fact  a  crystalline  substance,  HNO3-H2O,  which  is 
really  H3NO4,  is  actually  known  (541) .  For  this  hydrate  we  may 

write  the  formula 

HOv 

HO— N=O 
HO/ 
and  for  nitric  acid  itself 

//Q 

H0'<o 


556  Introduction  to  General  Chemistry 

Both  formulae  give  nitrogen  a  valence  of  five.     Orthophosphoric 
metaphosphoric  acids  probably  have  analogous  formulae. 
For  the  two  oxides  of  sulfur  we  write 


\) 

respectively,  thereby  assuming  the  valence  of  sulfur  to  be  four 
in  the  first  case  and  six  in  the  second.  The  corresponding  acids 
then  become 

HOV  HOV 

;>S  =  Oand         ;>S< 
HO/  HCK 

The  acids  of  the  seventh  group  of  the  type  HEO4  very 
probably  have  the  structure 

O 

II 
HO-E=O 

II 
O 

with  a  manganese  or  halogen  atom  of  valence  seven. 

Osmium,  the  only  element  of  the  eighth  group  which  forms 
an  oxide,  E04,  doubtless  has  a  valence  here  of  eight.  If  so  this 
oxide  must  be  represented  thus: 

O 

II 
0=Os=O 

II 
O 

840.  Positive  and  Negative  Valence. — We  have  already  seen 
that  the  valence  of  an  element  is  represented  by  the  number  of 
electrons  each  of  its  atoms  loses  or  gains  when  it  reacts  to  form 
a  compound  (484) .  Thus  when  copper  and  chlorine  unite  (246) 
to  form  CuCl2  each  atom  of  copper  loses  two  electrons  and  each 
atom  of  chlorine  gains  one.  When  sulfur  burns  to  form  S02 
(340)  each  atom  of  sulfur  loses  four  electrons  and  each  atom  of 
oxygen  gains  two.  But  when  sulfur  and  hydrogen  combine  to 
torm  H2S  it  is  the  hydrogen  which  loses  electrons,  since  this 
element  forms  positive  ions  only  (H+),  and  therefore  each  atom 


Classification  of  the  Elements.     The  Periodic  System      557 


of  sulfur  must  gain  two  electrons.  The  graphic  formulae  of  the 
two  sulfur  compounds  showing  the  distribution  of  charges 
resulting  from  the  transfers  of  electrons  are  as  follows: 


Sulfur  dioxide, 
Hydrogen  sulfide, 


H— S— H 


In  these  formulae  each  plus  sign  indicates  a  loss  of  one  electron. 
In  SO2  the  valence  of  sulfur  is  four;  in  H2S  it  is  two.  But  very 
plainly  these  are  different  kinds  of  valence,  and  we  should  dis- 
tinguish them  by  calling  the  first  a  positive  valence  and  the 
second  a  negative  valence. 

Other  elements  also  exhibit  both  positive  and  negative 
valence.  Thus  phosphorus  in  PH3  (588)  has  a  negative  valence 
of  three  and  in  PC13  (247,  576)  a  positive  valence  of  three,  while 
in  PC1S  it  has  a  positive  valence  of  five. 

TABLE  LII 


Group 

IV 

V 

VI 

VII 

VIII 

Compound  

CH4 

PH3 

H2S 

HC1 

Negative  valence  
Compound  
Positive  valence 

co2,4cci4 

4 

P2OS,  PC1S 

5 

S03,  H2S04 
6 

HC1O4 

o 
Os04 

8 

Snm  of  +and—  valence. 

8 

8 

8 

8 

8 

A  very  remarkable  law  can  be  formulated  from  facts  like 
those  brought  out  by  Table  LII.  For  many  elements  of  Groups 
IV  to  VIII  the  sum  of  the  maximum  positive  and  negative  valences 
of  an  element  is  eight. 

841.  Metals  and  Non-Metals. — We  may  now  consider  the 
way  in  which  the  metals  and  the  non-metals  are  distributed  in 
the  Periodic  Table  (829).  In  the  table  the  heavy  line  starting 
between  beryllium  and  boron  and  ending  with  the  boundary 
between  the  seventh  and  eighth  groups  separates  all  of  the 
pronounced  non-metallic  elements  from  the  others,  which  are 
all  more  or  less  metallic  in  character.  The  symbols  of  the  former 
are  printed  in  italics.  It  is  true,  however,  that  a  few  elements, 


558  Introduction  to  General  Chemistry 

particularly  chromium,  molybdenum,  and  manganese,  thus 
included  with  the  non-metals,  are  in  the  free  state  definitely 
metallic  and  form  salts  with  strong  acids,  e.g.,  Cr2(S04)3  and 
Mn(NO3)2.  But  it  is  also  a  fact  that  these  elements  all  form 
acids  that  yield  with  bases  typical  salts,  e.g.,  Na2Cr04,  Na2Mo04, 
and  KMnO4,  and  in  this  respect  the  elements  resemble  the  non- 
metals. 

It  is  particularly  the  non-metallic  elements  segregated  by 
the  heavy  line  of  Table  L  that  exhibit  both  positive  and  negative 
valence.  The  rare  gases  (Group  O)  have  little  in  common  with 
either  the  metallic  or  the  non-metallic  elements. 

842.  Chemical  Activity  and  the  Periodic  Table. — For  prac- 
tical purposes  we  may  define  the  chemical  activity  of  an  element 
as  its  tendency  to  enter  into  combination.     If  one  element 
displaces  another  from  the  solution  of  one  of  its  compounds 
(488-492)  the  first  is  the  more  active  of  the  two.     For  example, 
in  family  IA  the  displacement  order  is  Cu,  Ag,  Au,  and  this  is 
therefore  the  order  of  the  elements  in  respect  to  decreasing 
activity.     For  the  halogen  family,  VIIB,  the  displacement  order 
is  F,  Cl,  Br,  I,  and  this  again  is  the  order  of  decreasing  activity. 
A  thorough  study  of  the  displacement  tendencies  and  activities 
has    established    the   following   important   generalization:    In 
Groups  I  to  VII,  inclusive,  in  the  A  families  the  activity  of  an 
element  increases  with  increasing  atomic  weight;   in  the  B  series 
activity  decreases  as  the  atomic  weight  increases.     In  accord  with 
this  law  the  greatest  activities  are  found  in  the  elements  caesium, 
Cs,  at  the  bottom  of  IA  and  fluorine  at  the  top  of  VIIB.     It  is 
also  important  to  note  further  that  the  valence  of  caesium  in 
its  compounds  is  positive,  while  that  of  fluorine  is  negative. 

On  the  other  hand,  the  least  active  elements  are  met  with  at 
the  bottom  of  Group  VIII  in  the  metals  Os,  Ir,  and  Pt,  and  at 
the  top  of  Group  O,  where  we  find  helium  an  element  of  no 
chemical  activity  and  so  little  physical  affinity  (cohesion)  that 
in  liquid  form  it  has  the  lowest  boiling-point  of  all  the  elements 
and  is  the  only  element  that  has  not  been  solidified. 

843.  Periodic  Properties. — We  have  now  discussed,  as  far 
as  the  limits  of  this  text  will  permit,'  the  ways  in  which  the 


Classification  of  the  Elements.     The  Periodic  System      559 

properties  of  the  elements  change  as  we  pass  from  element  to 
element  in  each  family  and  from  family  to  family  and  group  to 
group  in  the  table.  It  now  remains  to  call  attention  to  the 
remarkable  manner  in  which  the  properties  of  the  elements 
change  as  we  pass  through  the  table  from  period  to  period  in  the 
order  of  increasing  atomic  weights.  The  periodic  fluctuation  of 
properties  is  most  easily  exhibited  by  means  of  a  diagram  or 
graph.  Let  us  take  for  illustration  the  variation  of  atomic 
volume  (823)  with  atomic  weight.  In  Fig.  116  the  atomic  vol- 
umes are  plotted  on  the  vertical  axis  and  the  atomic  weights 
on  the  horizontal  axis.  The  result  is  a  very  remarkable  graph, 
in  which  each  period  is  represented  by  a  crude  U-shaped 
curve.  We  see  that  there  is  a  periodic  repetition  of  like  properties 
as  we  pass  along  the  graph.  If  we  should  plot  in  a  similar  way 
any  one  of  many  other  properties  of  the  elements,  for  instance 
the  melting-points  or  the  compressibilities,  we  obtain  periodic 
curves  (graphs)  of  more  or  less  similar  forms.  These  facts  are 
summarized  by  the  statement  first  made  by  the  great  Russian 
chemist  Mendelejeff :  The  properties  of  the  elements  are  periodic 
functions1  of  their  atomic  weights.  This  means  that  the  atomic 
weight  of  an  element  determines  its  properties.  This  is  plain  if 
we  note  that  the  place  of  any  element  in  the  Periodic  Table  is 
fixed  by  its  atomic  weight,  and  that  the  properties  of  the  element 
are  indicated  by  its  position  in  the  table.  Among  the  properties 
of  elements  which  show  periodic  relations  we  may  mention 
valence,  chemical  activity,  melting-  and  boiling- temperatures, 
conductivity  for  heat  and  for  electricity,  hardness,  etc. 

844.  History  of  the  Periodic  System. — The  family  relation- 
ships of  the  elements  began  to  be  discovered  long  before  the 
Periodic  Table  as  a  whole  was  developed.  As  early  as  1829 
Dobreiner  pointed  out  the  existence  of  triads  of  elements,  such 
as  Cl,  Br,  I;  Li,  Na,  'K;  and  Ca,  Sr,  Ba,  in  which  the  atomic 
weight  and  other  properties  of  the  middle  element  of  the  triad 
were  close  to  the  mean  of  the  other  two.  As  time  went  on 
other  families  and  their  relations  were  recognized.  The  first 

1  The  term  function  is  much  used  in  mathematics.  If,  for  example,  the  value 
of  x  is  dependent  on  the  value  of  y,  then  x  is  said  to  be  a  function  of  y. 


Introduction  to  General  Chemistry 


attempt  to  arrange  all  of  the  elements  in  a  single  table  was  made 
by  Newlands  in  1864-66.     A  little  later  (1869-70)  Mendelejeff 


published  a  table  almost  in  the  form  of  Table  L.     There  were 
several  places  left  vacant  by  Mendelejeff  for  elements  as  yet 


Classification  of  the  Elements.     The  Periodic  System      561 

undiscovered,  and  the  whole  zero  group  was  of  course  absent, 
as  none  of  these  elements  was  discovered  until  many  years  later. 
It  was  Mendelejeff  who  pointed  out  many  of  the  relations 
between  the  properties  of  the  elements  in  essentially  the  form 
discussed  in  this  chapter  and  drew  the  far-reaching  conclusion 
that  the  properties  of  the  elements  are  periodic  functions  of  their 
atomic  weights.  This  statement  is  known  as  the  periodic  law. 
The  German  chemist  Lothar  Meyer  also  discovered  the  periodic 
law  independently  and  published  a  periodic  table  closely  resem- 
bling that  of  Mendelejeff  very  shortly  after  the  appearance  of 
the  latter's  table. 

845.  The  Vacant  Places  of  Mendelejeff's  Table.— In  Men- 
delejeff's  table  the  spaces  now  occupied  by  scandium,  gallium, 
and  germanium  were  left  vacant.  He  very  boldly  took  the  stand 
that  these  places  represented  elements  still  to  be  discovered.  In 
fact,  his  description  of  these  undiscovered  elements,  which  he 
called  ekaboron,  ekaluminum,  and  ekasilicon  respectively,  was 
so  complete  and  accurate  that  chemists  knew  just  how  these 
elements  would  behave  if  present  in  mixtures  with  other 
elements.  In  other  words,  they  knew  in  just  what  sorts  of 
chemical  residues  to  expect  these  new  elements,  and  it  was  not 
long  until  all  three  had  been  found.  In  1875  Lecoq  de  Bois- 
baudran  discovered  ekaluminum  and  called  it  gallium  (from 
Gaul,  in  honor  of  France);  in  1879  Nilson  discovered  ekaboron 
and  called  it  scandium  (for  Scandinavia) ;  and  in  1886  Winkler 

TABLE  LIII 

MENDELEJEFF'S  PREDICTIONS,  1871  WINKLER'S  OBSERVATIONS,  1886 

Ekasilicon;  £5=72  Germanium;   66=72.5 

Metal;  density  5. 5  Metal;  density  5. 47 

Oxide,  EsO2;  density  4.7  Oxide,  GeO2;  density  4. 7 

Chloride,  EsCl4;  density,  1.9  Chloride,  GeCl4;  density  i . 887 

Chloride  should  boil  below  100°        Chloride  boils  at  86° 
Fluoride,  EsF4;  not  gaseous  Fluoride,  GeF4;  solid 

Sulfide,  EsS2;  -insoluble  in  water,     Sulfide,  GeS2;  insoluble  in  water, 
but  soluble  in  ammonium  sulfide        but  soluble  in  ammonium  sulfide 
Tetra  ethyl  compound,  Es(C2Hs)4;    Tetra  ethyl  compound,  Ge(C2Hs)4; 
density  0.96;  boiling-point  160°       density  a  little  less  than  water; 

boiling-point  160° 


562  Introduction  to  General  Chemistry 

discovered  ekasilicon,  which  he  named  germanium  (for  Germany) . 
In  Table  LIII  some  of  the  properties  of  the  last-named  element 
are  set  down  opposite  those  predicted  by  Mendelejeff. 

Mendelejeff  also  predicted  the  existence  of  ekamanganese 
with  atomic  weight  about  100,  but  this  element  remains  as  yet 
unknown. 

When  radium  (480)  was  discovered  by  Mme  Curie  and  its 
atomic  weight  was  determined  to  be  226  it  was  at  once  placed  at 
the  bottom  of  IIA  as  a  member  of  the  alkaline  earth  family,  of 
which  its  chemical  properties  (see  chap,  xxxii)  show  it  to  be  a 
member.  The  zero  family  of  the  Periodic  Table  was  added  by 
Ramsay  after  the  discovery  of  the  inert  gases  (791-798)  and 
served  to  round  out  the  system  in  an  unexpected  but  most 
satisfactory  and  suggestive  fashion. 

846.  The  Rare  Earths  and  the  Periodic  Table.— The  fitting 
of  the  rare  earths  (820)  into  the  Periodic  Table  presents  a  diffi- 
culty that  is  not  yet  very  satisfactorily  solved.     These  ele- 
ments, as  we  have  seen,  are  all  trivalent  in  their  characteristic 
compounds  and  all  much  alike  in  their  properties.     They  plainly 
form  one  large  family  quite  as  homogeneous  as  any  other.     Only 
the  first  member,  lanthanum,  La,  and  possibly  also  the  second, 
cerium,  Ce,  fall  into  their  proper  places  when  these  elements  are 
arranged  in  the  order  of  their  increasing  atomic  weights.     The 
problem  of  the  position  of  the  rare  earths  in  the  table  is  an 
interesting  one  that  still  awaits  solution. 

847.  The  Position  of  Hydrogen  in  the  Periodic  Table. — As 
the  Periodic  Table  is  usually  written  hydrogen  is  not  included. 
The  chemistry  of  hydrogen  is  so  unique  that  this  element  is  in 
reality  in  a  class  by  itself.     Its  positive  valence  of  one  suggests 
a  place  in  the  first  group ;  but  solid  hydrogen  (296)  is  not  metallic 
and  bears  no  physical  resemblance  to  the  elements  of  Group  I. 
A  critical  discussion  of  the  position  of  hydrogen  in  the  periodic 
system  would  carry  us  beyond  the  scope  of  this  text. 

848.  Two  Striking  Anomalies. — In  Table  L  there  are  two 
striking  anomalies:   argon  with  an  atomic  weight  of  39.9  pre- 
cedes potassium  with  a  value  of  39.10;   and  tellurium,  127.5, 
precedes  iodine,  126.9.     Before  the  discovery  of  argon  the  case 


Classification  of  the  Elements.     The  Periodic  System      563 

of  tellurium  and  iodine  was  unique,  and  for  a  long  time  chemists 
were  inclined  to  think  that  the  then  accepted  values  of  the  atomic 
weights  of  these  two  elements  were  in  error,  for  it  seemed 
incredible  that  there  should  be  an  exception  to  a  law  that  held 
so  consistently  for  all  other  known  elements.  It  was,  of  course, 
out  of  the  question  to  place  iodine  before  tellurium,  for  this 
would  bring  iodine  in  VIB  with  sulfur  and  selenium  (815)  and 
place  tellurium  among  the  halogens!  It  was  possible  that 
either  or  both  of  the  elements  used  in  early  atomic-weight 
determinations  were  impure,  or  that  methods  of  analysis  were 
faulty,  and  that  in  consequence  the  atomic  weight  of  iodine 
should  be,  in  reality,  greater  than  that  of  tellurium.  The 
scientific  importance  of  the  problem  thus  presented  made 
chemists  realize  the  need  of  determining  atomic  weights  of  all 
elements  with  the  greatest  possible  accuracy.  As  a  result  some 
of  the  world's  most  skilled  chemists  have  given  their  best  efforts 
to  such  work.  In  the  cases  of  the  two  elements  under  discussion 
the  final  and  universally  accepted  results  are 

Tellurium  =127. 5 
Iodine       =126.92 

and  we  are  therefore  forced  to  conclude  that  the  position  of  an 
element  in  the  Periodic  Table  is  not  of  necessity  strictly  fixed  by  its 
atomic  weight.  The  more  recently  discovered  case  of  argon  and 
potassium  serves  to  confirm  this  conclusion. 

A  discovery  made  by  Mosely  in  1914  and  discussed  in  the 
following  chapter  has  thrown  a  riew  light  on  this  whole  matter. 
This  discovery  consists  in  nothing  less  than  the  finding  of  a 
property  of  the  elements  more  fundamental  than  their  atomic 
weights,  namely,  their  atomic  numbers.  If  the  elements  are 
arranged  according  to  these  numbers  the  two  anomalies  dis- 
appear and  every  other  element  falls  in  its  correct  position. 
But  this  is  another  story,  which  finds  its  logical  place  in  the 
following  chapter. 

849.  The  Newer  Tables:  The  Harkins  Table. — Since  the 
time  of  Mendelejeff  many  other  tables  have  been  arranged. 
We  shall  have  space  to  present  only  one  of  these,  namely,  that 


564  Introduction  to  General  Chemistry 

of  Professor  W.  D.  Harkins.  This  table  should  be  of  particular 
interest  to  those  who  expect  to  specialize  in  chemistry,  since  it 
shows  the  relationships  between  the  elements  much  more  clearly 
than  does  the  Mendelejeff  table.  Fig.  117  is  a  projection  of  a 
space  model  of  the  table. 

The  elements  are  arranged  about  a  continuous  helix  in  such 
a  way  that  those  of  the  same  family  are  in  the  same  vertical 
column.  In  the  actual  model  the  path  of  the  helix  is  indicated 
by  a  heavy  wire  connecting  balls  which  are  lettered  with  the 
symbols  of  the  elements  which  they  represent.  The  families 
(831)  are  also  connected  by  heavy  vertical  rods  of  brass,  which 
thus  form  the  supporting  structure  of  the  model.  Each  unit 
of  atomic  weight  is  represented  by  i  cm.  vertical  distance  on  the 
spiral  (see  scale  at  the  side  of  Fig.  117).  The  elements  are 
arranged  in  the  order  of  their  atomic  numbers. 

The  helix  begins  with  hydrogen,  which  is  uniquely  placed 
at  the  top  of  the  table.  Next  the  helix  descends  to  the  posi- 
tion of  helium  and  makes  a  big  loop,  coming  back  to  neon,  which 
is  of  course  just  under  helium.  On  this  loop  are  found  the 
elements  of  the  first  period  of  the  Mendelejeff  table,  namely 
lithium,  beryllium,  boron,  carbon,  nitrogen,  oxygen,  and  fluorine. 
After  arriving  at  the  position  of  neon  the  helix  makes  another 
big  loop  back  to  argon,  and  on  this  are  the  elements  of  the 
second  period  of  the  Mendelejeff  table.  These  are  sodium, 
magnesium,  aluminum,  silicon,  phosphorus,  sulfur,  and  chlorine. 
From  argon  the  helix  passes  next  to  potassium,  calcium,  scan- 
dium, and  titanium;  then  it  begins  the  first  inner  loop,  passing 
successively  through  the  positions  of  vanadium,  chromium, 
manganese,  iron,  cobalt,  nickel,  copper,  zinc,  and  gallium  before 
returning  to  the  outer  loop  at  the  position  of  germanium.  Next 
the  helix  takes  a  full  outer  turn  through  the  positions  of  arsenic, 
selenium,  bromine,  krypton,  etc. 

If  we  continue  following  the  helix  we  find  that  the  first  four 
A-family  elements  and  the  last  four  B -family  elements  are  on 
the  outer  loops,  together  with  the  zero  group.  The  second  four 
A-family  elements,  the  eighth  group,  and  the  first  four  B-family 
elements  are  on  the  inner  loops.  This  brings  corresponding  A 


Classification  of  the  Elements.     The  Periodic  System     565 


Periodic  Table   by  WD.Harkins 
FIG.  IT? 


566  Introduction  to  General  Chemistry 

and  B  families  together  in  a  very  ingenious  way.  At  the  bottom 
of  the  model  a  dotted  line  is  traced  to  show  the  relative  positions 
of  the  inner  and  outer  loops.  It  will  be  noticed  that  the  inner 
loops  spring  from  the  outer  at  the  fourth  group,  and  that  from 
that  point  the  distance  between  the  two  increases,  becoming 
greatest  between  the  zero  and  the  eighth  group.  The  rare 
earths  are  placed  on  the  vertical  column  of  IIIA,  since  their 
chemical  properties  show  them  to  be  closely  related  to  this 
family.  The  last  loop  of  the  helix  represents  the  positions 
of  the  radioactive  elements,  which  will  be  discussed  in  the 
next  chapter. 

This  table  has  the  advantage  of  representing  the  elements  in 
a  continuous  series.  It  also  shows  the  relationship  of  the  A  and 
B  families  more  plainly  than  does  the  Mendelejeff  table.  The 
vacancies  in  this  table  are  for  atomic  numbers  as  yet  unassigned 
to  any  element,  and  there  is  therefore  a  very  strong  probability 
that  elements  will  be  found  to  occupy  these  places.  This  was  of 
course  not  true  of  many  of  the  vacant  spaces  of  the  Mendelejeff 
table.  Professor  Harkins  counts  hydrogen  and  helium  as  the 
first  period  and  calls  it  the  zero  period.  The  subsequent  periods 
he  counts  as  beginning  with  the  alkalies  and  ending  with  the 
noble  gases.  Thus,  period  one  would  begin  with  lithium  and 
end  with  the  neon. 

850.  Summary. — In  this  chapter  we  have  discussed  a  system 
of  classification  which  brings  all  the  elements  into  a  single  table 
(Table  L,  829),  in  which  the  individuals  are  arranged  in  nine 
groups  (vertical  columns)  and  seven  periods  (horizontal  rows). 
The  first  two  periods  are  short  ones,  the  others  long  (829) ;  and 
in  the  latter  the  elements  of  Group  I  to  VII  are  further  classified 
into  A  and  B  subgroups  or  families  (83 1 ) .  In  the  first  two  (short) 
periods  the  elements  of  Groups  I  and  II  belong  to  the  correspond- 
ing A  families  (831,  832),  while  in  the  same  periods  the  members 
of  Groups  VI  and  VII  are  B-family  elements  (836,  837).  The 
first-  and  second-period  members  of  Groups  III,  IV,  and  V  may 
be  classed  with  either  the  respective  A  or  B  families. 

In  the  periodic  system  of  arrangement  most  of  the  properties 
of  elements  and  their  compounds  vary  gradually  as  we  pass 


Classification  of  the  Elements.     The  Periodic  System      567 

from  one  element  to  its  neighbor  in  the  table.  This  is  true 
whether  we  pass  from  top  to  bottom  of  the  vertical  columns, 
considering  each  family  separately,  or  pass  through  any  one  of 
the  seven  periods  in  the  order  of  increasing  atomic  weights. 
Within  a  given  family  the  properties  change  in  a  systematic 
manner  in  such  a  way  that  a  knowledge  of  the  chemical  properties 
and  the  behavior  of  a  given  member  can  be  deduced  from  those 
of  each  of  its  nearest  relations.  This  is  one  of  the  most  impor- 
tant facts  about  the  periodic  system.  Furthermore  in  each 
period  we  find  a  recurrence,  in  like  order  but  of  different  degree, 
of  the  properties  that  appear  in  all  the  periods.  In  consequence, 
if  we  plot  the  graph  of  any  given  property,  taking  the  elements 
in  serial  order  (in  the  order  of  increasing  atomic  numbers,  848), 
we  get  a  periodic  curve  more  or  less  closely  resembling  in  form 
that  for  atomic  volumes  shown  in  Fig.  116.  These  facts  are 
epitomized  in  the  periodic  law  of  Mendelejeff:  "  The  properties 
of  the  elements  are  periodic  functions  of  their  atomic  weights" 

The  simple  and  systematic  way  in  which  valence  (both  posi- 
tive and  negative)  changes  from  group  to  group  (839,  840)  is  one 
of  the  most  striking  and  useful  facts  brought  out  in  this  chapter, 
and  one  that  should  be  firmly  fixed  in  the  mind  of  every  chemist. 

Finally  the  Periodic  Table  is  of  great  value  to  all  students  of 
chemistry  in  presenting  to  them  a  bird's-eye  view  of  all  the 
elements.  In  so  doing  it  furnishes  a  mass  of  evidence  for  the 
unity  of  matter  in  the  sense  indicated  earlier  (470,  482),  as  will 
be  discussed  more  fully  in  the  following  chapter.  It  seems  to 
set  a  limit  to  the  possible  number  of  elements  and  also  to  indicate 
the  nature  of  the  yet  undiscovered  ones.  Of  more  practical 
importance  than  all  else  the  table  makes  possible  a  system  of 
classification  for  an  enormous  number  of  chemical  facts  and  thus 
permits  through  its  mastery  an  orderly  arrangement  of  the 
knowledge  that  constitutes  the  science  of  chemistry. 


CHAPTER  XXXII 
RADIOACTIVITY  AND  THE  NATURE  OF  MATTER 

851.  Introduction. — The  study  of  the  history  of  chemistry 
shows  very  plainly  that  the  rapid  growth  of  chemical  knowledge 
was  in  no  small  measure  due  to  the  innumerable  attempts  of 
the  alchemists  to  convert  the  base  metals  into  geld.  Before  the 
dawn  of  scientific  chemistry  belief  in  the  possibility  of  trans- 
mutation of  the  elements  was  nearly  universal.  As  time  went 
on  and  the  facts  and  ideas  discussed  in  the  early  chapters  of  this 
book  became  known,  alchemy  merged  gradually  into  chemistry, 
which  soon  became  a  science  as  well  as  an  art.  With  the  science 
of  chemistry  came  the  conviction  that  the  elements  were 
immutable  kinds  of  matter,  and  that  transmutation  was  a 
myth  consisting  of  nothing  more  substantial  than  the  hopes  of 
the  avaricious  or  the  cupidity  of  charlatans. 

After  some  years  of  disrepute  new  life  was  injected  into  the 
old  idea  in  1814  by  a  suggestion  of  Prout.  This  Englishman 
called  attention  to  the  fact  that  the  atomic  weights  of  the 
gaseous  elements  were  whole  numbers  and  therefore  even  mul- 
tiples of  the  atomic  weight  of  hydrogen  and  expressed  the  opinion 
that  the  atoms  of  other  elements  were  made  up  of  atoms  of 
hydrogen.  This  idea  came  to  be  known  as  Prout 's  hypothesis. 
The  first  defeat  for  the  hypothesis  came  when  it  was  proved  that 
the  atomic  weight  of  chlorine  was  35 .5.  Prout  now  claimed  a 
half-atom  of  hydrogen  as  the  unit.  Later  work  on  atomic 
weights  (see  Table  XLI,  800)  proved  even  this  half-size  unit  too 
large  if  the  hypotheses  were  to  be  applied  to  all  elements.  On 
the  whole  Prout  had  but  few  supporters.  Nevertheless  it  is  an 
incontestable  fact  that  for  many  elements  the  atomic  weights 
are  much  nearer  whole  numbers  than  can  be  accounted  for  by 
chance. 

The  discovery  of  the  periodic  law  furnished  new  and  to 
many  the  most  convincing  evidence  that  all  elements  are  but 

568 


Radioactivity  and  the  Nature  of  Matter  569 

modifications  of  one  primitive  form  of  matter.  But  no  satis- 
factory theory  of  the  nature  of  matter  and  the  interrelations  of 
elements  was  developed  until  the  knowledge  of  the  main  facts 
of  radioactivity  (480)  had  led  to  the  disintegration  hypothesis 
as  their  most  plausible  explanation.  Since  the  discovery  of 
radium  and  radioactive  phenomena  has  furnished  the  key  to 
one  of  nature's  most  profound  mysteries  we  shall  devote  the 
present  chapter  to  an  account  of  this  subject.  The  story  of 
the  discovery  of  radium  is  most  instructive  in  illustrating  the 
way  in  which  scientific  advances  of  the  greatest  significance  often 
result  from  the  thorough  investigation  of  phenomena  that  seem 
trivial  to  persons  who  call  themselves  practical.  We  shall  there- 
fore relate  the  facts  in  their  historical  sequence. 

852.  The  X-Rays  and  Phosphorescence. — The  X-ray  appa- 
ratus has  already  been  described  (476),  and  it  has  been  men- 
tioned that  the  X-rays  are  produced  when  the  cathode  rays 
strike  the  target.    It  has  also  been  stated  that  the  cathode 
rays,  non-luminous  themselves,  cause  the  glass  walls  of  a  cathode- 
ray  tube  to  glow  with  a  greenish-yellow  phosphorescence.    Thus 
there  seemed  to  be  a  causal  relation  between  phosphorescence 
and  X-rays.     This  supposed  relationship  was  made  the  subject 
of  an  extensive  and  thorough  investigation  by  Henri  Becquerel, 
professor  of  physics  at  the  Sorbonne  (University  of  Paris).     It 
had  long  been  known  that  a  number  of  substances  gave  out  a 
faint  phosphorescent  light  for  some  time  after  exposure  to  any 
bright  light.    All  such  phosphorescent  substances  were  carefully 
examined  by  Becquerel  to  learn  if  these  glowing  bodies  gave  out 
X-rays.    The  latter  have  the  same  effect  on  photographic  plates 
that  light  has,  but  they  can  pass  directly  through  lightproof 
paper.    Accordingly  Becquerel  tested  the  effect  of  these  phos- 
phorescent substances  on  plates  wrapped  in  black  paper.    None 
showed  any  indication  of  the  production  of  X-rays,  with  one 
exception:  this  was  potassium  uranyl  sulfate,K2U02(SO4)2  •  4H2O, 
a  salt  of  the  rather  uncommon  element  uranium  (836). 

853.  The    Becquerel    Rays. — Not    only    did    the    above- 
mentioned  uranium  salt  give  out  rays  that,  like  X-rays,  pene- 
trated black,  lightproof  paper  and  acted  on   a  photographic 


570  Introduction  to  General  Chemistry 

plate  while  phosphorescing  after  it  had  been  exposed  to  sun- 
light, but  the  salt  showed  equal  photoactivity  if  prepared  and  also 
tested  wholly  in  the  dark,  under  which  condition  it  did  not  phos- 
phoresce. Furthermore  all  other  uranium  compounds  tested 
showed  similar  photoactivity ,  although  they  were  not  phos- 
phorescent even  if  previously  exposed  to  light.  Obviously 
these  new  rays,  called  at  first  Becquerel  rays,  were  not  caused 
by  phosphorescence  but  were  produced  by  the  element  uranium 
in  all  forms  of  chemical  combination. 

854.  lonization  of  Air  by  Becquerel  Rays. — It  was  also  soon 
discovered  that  Becquerel  rays  ionize  air  (770),  so  that  by 
means  of  the  discharge  of  a  gold-leaf  electroscope  the  activity 
of  a  substance  could  be  tested  much  more  readily  than  by 
the  photographic  method.    A  convenient  form  of  electroscope 
suitable  for  projection-lantern  experiments  is  shown  in  Fig.  109 
(770).    Two  parallel  sides  are  made  of  glass  plates;  the  rest  of 
the  case  is  of  metal.    A  wire  passes  through  the  amber  plug 
insulator  and  ends  in  a  brass  strip  6  to  8  mm.  wide  and  5  or  6  cm. 
long.    A  single  leaf  of  gold  or  aluminum  is  attached  to  the  brass 
strip,  as  shown  in  the  figure.    The  electroscope  may  be  charged 
by  means  of  an  ebonite  rod,  celluloid  comb,  fountain-pen  holder, 
etc.,  electrified  by  being  rubbed  with  a  piece  of  woolen  cloth. 
The  uranium  compound  is  held  within  2  or  3  cm.  of  the  end  of 
the  wire  carrying  the  gold-leaf  system.     The  rate  of  discharge 
of  the  electroscope  under  fixed  conditions  is  a  measure  of  the 
intensity  of  the  Becquerel  rays. 

855.  The    Discoveries    of    Mme    Curie. — BecquerePs    dis- 
covery of  the  rays  known  by  his  name  was  published  in  1896, 
about  two  years  after  Roentgen  had  discovered  X-rays.    At 
about  this  time  Becquerel  suggested  to  Mme  Curie,  the  brilliant 
wife  of  Becquerel's  colleague,  Professor  Curie,  that  the  investiga- 
tion of  natural  uranium  minerals  should  prove  interesting. 
Mme  Curie's  experiments  soon  showed  that  all  uranium  minerals 
were  photographically  active;    they  were  considerably  more 
active  in  ionizing  air  than  one  would  expect  from  their  uranium 
content.     This  fact  seemed  to  suggest  that  the  larger  part  of 
the  activity  of  the  minerals  was  eliminated  during  the  chemical 


Radioactivity  and  the  Nature  of  Matter  571 

treatment  used  in  making  the  pure  uranium  compounds.  Search 
was  therefore  made  for  other  active  substances  among  the 
by-products  of  the  preparation  of  uranium  salts;  whereupon  it 
was  discovered  that  the  small  amount  of  barium  extracted  from 
the  waste  materials  was  many  times  as  active  as  any  pure 
uranium  salt,  notwithstanding  the  fact  that  ordinary  barium 
salts  were  entirely  inactive. 

856.  The  Discovery  of  Radium.  —After  considerable  quanti- 
ties of  active  barium  chloride  had  been  extracted  from  uranium 
mineral  residues  Mme  Curie  undertook  to  separate  from  the 
barium  the  new  substance  to  which  she  attributed  the  activity. 
She  found  that  if  a  hot,  saturated  solution  of  the  active  barium 
chloride  was  allowed  to  cool  about  half  of  the  dissolved  salt 
crystallized  out,  but  these  crystals  contained  70  per  cent  or 
more  of  the  active  material.  Thus  the  crystals  were  appre- 
ciably richer  in  the  active  matter  than  the  original  material. 
By  repeated  recrystallization  of  the  crystals  that  separated  in  this 
way  Mme  Curie  finally  obtained  a  product  that  was  thousands 
of  times  more  active  than  an  equal  weight  of  a  uranium  salt. 
This  product  was  still  largely  barium  chloride;  but  that  it  also 
contained  a  new  element  was  shown  conclusively  by  the  fact 
that  its  spectrum  (786)  contained  in  addition  to  those  of  barium 
several  lines  not  belonging  to  any  known  element.  It  was 
therefore  certain  that  a  new  element  had  been  discovered,  and 
it  seemed  probable  that  it  was  this  element  which  had  produced 
the  powerful  Becquerel  rays  that  caused  the  photoactivity  and 
ionizing  power  of  the  material.  On  account  of  its  ability  to 
send  out  rays  or  radiations  Mme  Curie  named  the  supposed 
new  element  radium. 

The  next  problem  was  to  free  the  radium  salt  from  the  accom- 
panying barium.  This  was  finally  accomplished  by  a  long  series 
of  crystallizations  of  the  chlorides  of  the  barium-radium  mixture. 
Finally  pure  radium  chloride  free  from  barium  was  obtained.  It 
was  a  white  crystalline  salt  closely  resembling  barium  chloride 
in  appearance.  Its  activity,  called  now  radioactivity,  was  two 
or  three  million  times  as  great  as  that  of  an  equal  weight  of  a 
uranium  salt. 


572  Introduction  to  General  Chemistry 

857.  The  Atomic  Weight  of  Radium  and  Its  Position  in  the 
Periodic  Table. — If  an  element  does  not  form  volatile  com- 
pounds its  atomic  weight  cannot  be  determined  by  means  of  the 
law  of  minimum  weights  (64) ;  and  if  it  is  not  obtainable  in  con- 
siderable quantity  in  uncombined  solid  form  the  method  based 
on  the  law  of  Dulong  and  Petit  (229)  is  also  unavailable.  One 
might  make  use  of  a  method  based  on  direct  or  indirect  deter- 
mination of  the  osmotic  pressure  of  a  solution  (717-722),  but 
this  involves  a  complication  in  case  an  ionized  salt  is  used.  As 
none  of  the  ways  just  mentioned  could  be  applied  in  the  case  of 
radium,  of  which  but  a  fraction  of  a  gram  was  available  and  that 
worth  thousands  of  dollars  in  its  cost  of  preparation,  other  means 
of  fixing  its  atomic  weight  were  used. 

If  we  analyze  a  chloride  of  a  new  element  and  find  its  per- 
centage composition,  one  additional  fact  must  be  known  in 
order  to  fix  the  atomic  weight  of  the  element,  namely,  its  valence. 
But  if  we  can  discover  to  what- group  of  the  Periodic  Table  the 
new  element  belongs  we  can  of  course  infer  its  valence  (839). 
We  have  seen  that  radium  resembles  barium  so  closely,  in  some 
respects  at  least,  that  the  separation  of  the  chlorides  of  the  two 
elements  is  very  difficult.  This  indicates  that  radium  is  a  mem- 
ber of  the  alkali  earth  group  containing  calcium,  strontium,  and 
barium  (824,  832).  Furthermore  it  was  found  that  radium  also 
resembles  barium  in  the  following  respects:  its  carbonate  is 
insoluble  in  water,  its  sulfate  is  insoluble  in  water  and  in  strong 
acids,  and  its  hydroxide  is  soluble,  since  radium  chloride  solution 
gives  no  precipitate  with  either  ammonia  or  sodium  hydroxide. 
Furthermore  the  free  element  cannot  be  deposited  on  a  platinum 
electrode  from  a  water  solution  by  electrolysis.  It  therefore 
seemed  very  probable  that  radium  was  a  second-group  element 
and  had  a  valence  of  two.  The  analysis  of  the  pure  chloride 
was  made  by  Mme  Curie;  the  result,  with  an  assumed  valence  of 
two,  led  to  a  calculated  atomic  weight  of  226.  Since  the  space 
in  the  Periodic  Table  (829)  corresponding  to  this  atomic  weight 
was  vacant,  no  doubt  remained  of  the  correctness  of  the  atomic 
weight  assigned  to  the  new  element. 


Radioactivity  and  the  Nature  of  Matter  573 

858.  The  Alpha  Rays. — It  has  already  been  stated  that 
radium  emits  three  kinds  of  rays,  designated  as  alpha,  beta,  and 
gamma,  although,  strictly  speaking,  only  the  last  named  is  a  true 
radiation  (that  is,  a  vibration  in  the  luminous  ether),  since  the 
alpha  rays  are  atoms  of  helium  (795)  and  the  beta  rays  electrons 
(480).  Entirely  reliable  experiments,  which  are  too  intricate 
to  be  described  here,  have  shown  that  the  alpha  particles,  as  we 
shall  now  call  them,  are  shot  out  from  radioactive  atoms  with 
velocities  of  5,000  to  12,000  miles  a  second.  They  travel  in 
straight  lines  through  air  for  distances  (called  their  ranges)  of 
3  to  8  cm.  before  they  slow  down  to  the  average  velocity  of  air 
molecules  (about  one-fourth  of  a  mile  a  second,  197).  The 
alpha  particle  of  radium  has  a  range  of  3.5  cm.,  and  in  this 
distance  it  collides  with  many  thousands  of  air  molecules;  but 
instead  of  rebounding  as  an  elastic  body  and  changing  its  direc- 
tion after  each  collision  the  helium  atom  constituting  an  alpha 
particle  passes  straight  through  the  air  molecule,  as  a  rule,  with- 
out changing  its  direction  of  motion.  This  fact  furnishes  the 
most  convincing  evidence  that  molecules  are  not  solid  but  are 
made  up  of  a  group  of  widely  spaced  particles,  the  electrons  (470, 
483).  It  is  the  alpha  particles  which  cause  the  major  part  of 
the  ionization  of  air  exposed  to  radium  rays.  Each  alpha  particle 
leaves  in  its  wake  thousands  of  ions:  each  ion  is  an  air  molecule 
from  which  an  electron  has  been  detached,  leaving  a  positive  ion ; 
or  to  which  a  detached  electron  has  united,  forming  a  negative 
ion.  All  air  ions  carry  single  electric  charges  only.  Alpha 
particles  are  easily  stopped  by  thin  sheets  of  any  solid  substance, 
such  as  paper,  glass,  or  metal. 

Proof  that  the  alpha  rays  are  helium  atoms  was  conclusively 
furnished  in  the  following  way:  Radium,  contained  in  a  sealed 
glass  tube  with  walls  so  thin  that  the  alpha  particles  could  barely 
pass  through,  was  surrounded  by  a  sheet  of  lead  which  completely 
stopped  the  particles.  After  several  hours  the  lead  was  removed 
and  sealed  in  a  larger  glass  tube  having  wires  for  the  passage  of 
electric  sparks.  The  lead  was  then  melted  by  heating  the  glass 
tube;  this  set  free  the  helium  atoms  that  were  imbedded  in  the 


574 


Introduction  to  General  Chemistry 


lead,  as  was  proved  by  the  fact  that  on  passing  an  electric  dis- 
charge through  the  tube  the  spectrum  of  helium  was  easily 
recognized  by  observation  with  a  spectroscope  (785). 

859.  Luminous  Effect  of  the  Alpha  Rays. — When  alpha  rays 
strike  a  screen  covered  with  a  layer  of  a  specially  prepared  form 
of  crystalline  zinc  sulfide,  ZnS,  a  pale  phosphorescent  light  is 
produced.  Luminous  watch  dials,  which  have  recently  become 
so  popular,  are  produced  on  this  principle.  From  one  to  ten 
parts  of  radium  to  ten  thousand  parts  of  zinc  sulfide  are  used  in 
making  the  luminous  paint  for  such  purposes.  Radium  luminous 


FIG.  118 

paint  was  extensively  used  in  the  war  for  making  luminous  dials 
for  aeroplane  instruments,  submarine  instruments,  trench  com- 
passes, gun  sights,  etc. 

If  one  views  in  the  dark,  by  means  of  a  good  lens,  a  zinc 
sulfide  screen  exposed  to  alpha  rays  from  a  minute  amount  of 
radium  a  very  beautiful  effect  will  be  seen.  Instead  of  a  uniform 
glow  sparks  of  light  will  be  seen  appearing  here  and.  there 
at  random  over  the  screen.  It  is  now  certain  that  each  tiny 
flash  of  light  is  the  result  of  the  striking  of  a  single  helium  atom. 
The  simple  device  consisting  of  a  zinc  sulfide  screen,  a  bit  of 
radium,  and  a  lens  was  called  a  spinthariscope  by  Crookes  (477), 
its  inventor.  The  same  effect  may  be  seen  if  the  radium  paint 


Radioactivity  and  the  Nature  of  Matter 


575 


on  a  luminous  watch  dial  is  examined  in  the  dark  with  a  powerful 
lens. 

860.  The  Beta  Rays. — The  beta  rays  are  electrons  (480)  shot 
out  from  radioactive  atoms  with  velocities  closely  approaching 
the  velocity  of  light.  The  beta  rays  are  much  more  penetrating 
than  the  alpha  rays  and  can  pass  through  ten  to  twenty  sheets 
of  paper.  It  is  chiefly  these  rays  (and  in  small  measure  the 
gamma  rays)  that  are  photographically  active.  Fig.  118  shows 
a  photograph  taken  by  radium  rays.  The  photographic  plate 
was  wrapped  in  lightproof  paper;  the  objects  shown  in  the 
figure  were  laid  on  the  paper  and  the  whole  exposed  to  the 
radium  at  a  distance  of  10  cm.  for  a  few  minutes. 
The  plate  was  then  developed  in  the  usual  way,  and 
from  the  negative  so  obtained  the  positive  print 
reproduced  in  the  figure  was  made. 

A  curious  radium  clock  has  been  constructed  by 
R.  J.  Strutt,  an  English  physicist,  to  show  the  con- 
tinuous loss  of  negative  electricity  as  electrons  by 
radium.  The  radium  is  inclosed  in  a  small  tube,  A , 
Fig.  119,  made  of  glass  thick  enough  to  prevent  the 
alpha  rays  from  escaping  but  not  thick  enough  to 
stop  the  beta  rays.  This  small  tube  is  attached  at 
the  top  to  an  insulating  rod,  B,  of  quartz  and  carries 
at  its  lower  end  a  pair  of  gold  leaves,  C.  The  tube 
and  gold  leaves  are  fixed  in  a  glass  vessel  having  a 
pair  of  tin-foil  strips,  D,  pasted  on  the  inside  opposite  the  gold 
leaves.  These  strips  are  electrically  connected  to  a  wire  sealed 
through  the  glass  and  intended  to  furnish  a  connection  to  the 
earth.  The  larger  glass  vessel  is  highly  evacuated.  The  action 
of  the  apparatus  is  as  follows:  The  beta  rays  (electrons)  from  the 
radium  tube  escape  to  earth  by  way  of  the  tin-foil  strips,  D. 
This  leaves  the  radium  tube  with  a  deficiency  of  negative  elec- 
tricity and  therefore  positively  charged.  The  positive  charge 
causes  the  gold  leaves  to  diverge  (by  reason  of  the  repulsion  of 
like  charges)  until  they  touch  the  tin-foil  strips,  when  they  are 
at  once  discharged  and  fall  together.  Then  the  whole  cycle 
begins  again.  Since  the  discharge  of  the  gold  leaves  occurs  at 


FIG.  119 


576  Introduction  to  General  Chemistry 

regular  intervals  Strutt  called  the  device  a  radium  clock.  If  the 
vessel  were  not  evacuated  ionization  of  the  air  would  occur  and 
conduct  electricity  away  so  fast  that  the  gold  leaves  would  not 
become  charged.  The  beta  rays  can  also  ionize  gases  but  are 
less  effective  than  the  alpha  rays. 

861.  The  Gamma  Rays  and    X-Rays. — The   gamma  rays 
closely  resemble  X-rays.     They  are  far  more  penetrating  than 
the  beta  rays  and  can  pass  through  several  millimeters  of  metal 
with  little  absorption  and  are  only  completely  absorbed  by  several 
centimeters  of  lead,  which  is  more  effective  in  stopping  these 
rays  than  any  other  common  substance.     The  gamma  rays 
ionize  gases  and  also  affect  a  photographic  plate.     Several 
minerals  glow  perceptibly  when  exposed  to  the  gamma  rays  of 
a  considerable  quantity  of  radium  (say  loomg.).     Willemite,  a 
silicate  of  zinc  (807) ,  is  especially  active  in  this  way.     Diamonds 
also  glow  with  a  clear  white  light  under  the  gamma  rays,  so  that 
these  rays  may  be  used  practically  to  distinguish  a  real  from  an 
imitation  diamond. 

The  true  physical  nature  of  gamma  rays  and  X-rays  was  for 
a  long  time  in  doubt.  The  critical  work  which  cleared  up  the 
question  will  be  discussed  later  (877) . 

862.  Heat  from  Radium. — Not  long  after  radium  had  been 
obtained  in  considerable  quantity  it  was  discovered  that  a  tube 
of  this  unique  element  was  somewhat  warmer  than  its  surround- 
ings.    Subsequent  investigation  by  Rutherford  (481)   showed 
that  one  gram  of  radium  gives  out  134  calories  per  hour  and 
keeps  up  this  rate  of  heat  production  continuously  with  appar- 
ently little  change. 

863.  Radium  Emanation. — The  ionization  of  a  gas  does  not 
occur  spontaneously,  as  does  the  ionization  of  an  electrolyte. 
Moreover,  very  quickly  after  the  ionizing  agent  (say  a  tube  of 
radium)  is  removed  the  ions  of  a  gas  neutralize  one  another's 
charges  and  the  gas  becomes  nonconducting.     With  this  fact 
in  mind  let  us  consider  the  following  curious  observation:   If  a 
current  of  air  is  led  over  a  radium  compound,  or  better  through 
a  solution  of  a  radium  salt,  the  air  will  become  strongly  ionized 
and  remain  so  for  hours  and  even  days,  losing  its  electrical  con- 


Radioactivity  and  the  Nature  of  Matter  577 

ductivity  so  slowly  that  only  after  about  four  days  is  half  of 
the  effect  lost.  In  fact,  it  seemed  as  if  the  air  had  taken  up 
radioactive  matter  from  the  radium,  since  this  active  air  also 
caused  a  zinc  sulfide  screen  to  glow  (859),  and  this  glow  was 
found  by  examination  with  a  lens  to  be  the  result  of  scintillating 
flashes  like  those  caused  by  alpha  rays.  In  consequence  the 
radium  was  said  to  have  given  off  a  radioactive  emanation  to  the 
air.  The  material  nature  of  the  emanation  was  first  clearly 
shown  by  experiments  by  Rutherford,  in  which  air  containing 
emanation  was  passed  through  a  metal  tube  cooled  externally 
by  liquid  air  (776).  Every  trace  of  emanation  was  removed 
from  the  air  that  passed  through  the  cooled  tube ;  but  when  the 
latter  was  removed  from  the  liquid  air  and  allowed  to  come 
to  room  temperature  the  whole  of  the  emanation  was  recovered 
by  blowing  fresh  air  through  the  tube.  The  explanation  of 
this  experiment  given  by  Rutherford  is  very  simple:  Radium 
emanation  is  a  gaseous  radioactive  substance  which  condenses 
practically  completely  at  the  temperature  of  liquid  air.  Upon 
becoming  warm  it  again  volatilizes.  This  explanation  has 
proved  to  be  correct  in  every  particular. 

864.  The  Formation  and  Decay  of  Radium  Emanation. — The 
whole  of  the  emanation  contained  in  a  solution  of  a  radium  salt 
can  be  removed  by  blowing  a  current  of  air  through  the  solution 
for  a  few  minutes.  If  the  solution  is  sealed  air-tight,  new  emana- 
tion will  gradually  accumulate  in  it.  After  about  40  days  there 
is  present  in  the  sealed  vessel  a  maximum  of  emanation  strictly 
proportional  to  the  quantity  of  radium  present.  The  emanation 
accumulates  at  a  regular  rate,  as  follows:  after  3.85  days 
from  the  time  the  solution  was  sealed  (after  all  emanation  has 
previously  been  removed)  half  of  the  maximum  quantity  of 
emanation  is  present;  after  2X3.85  days,  three-fourths  of  the 
maximum;  after  3X3  .85  days,  seven-eighths  of  the  maximum, 
etc.,  so  that  after  40  days  over  99 .9  per  cent  of  the  maximum  is 
present.  We  have  seen  that  half  of  the  activity  of  a  given 
portion  of  emanation  is  lost  in  about  4  days  (863).  Strictly 
speaking,  exactly  half  is  lost  in  3  .85  days;  in  2X3  .85  days  half 
of  the  remainder  is  lost,  or  a  total  of  three-fourths;  in  3X3  -85 


Introduction  to  General  Chemistry 


days  half  of  the  remainder  is  again  lost,  or  a  total  of  seven- 
eighths,  etc.  The  emanation  is  said  to  decay  to  half  value  in 
3 . #5  days.  The  interval  of  3 . 85  days  is  called  the  period  of  half 
decay,  or  briefly  the  period.  The  rates  of  formation  and  decay 
of  radium  emanation  are  shown  by  the  two  graphs  of  Fig.  120. 
The  proportion  of  emanation  contained  in  a  sealed  sample  of 
radium  after  an  interval  of  40  days  is  a  maximum,  because  from 

this  time  on  the  decay  of  the 
emanation  present  exactly 
compensates  the  formation 
of  new  emanation. 

865.  Theory  of  Radio- 
active Change. — The  many 
facts  already  cited  show 
that  radium  has  properties 
that  place  it  in  a  class  by 
itself.  For  a  number  of 
years  after  the  discovery  of 
radium  the  cause  of  its 

strange  behavior  was  the  most  conspicuous  of  all  scientific 
mysteries.  Some  faint-hearted  scientists  declared  that  the 
known  facts  disproved  the  two  most  fundamental  scientific 
laws:  the  law  of  the  conservation  of  matter  and  the  law  of  the 
conservation  of  energy.  It  was  also  said  that  a  serious  incon- 
sistency existed  in  the  statement  that  radium  is  an  element,  if 
it  is  really  true  that  radium  gives  rise  to  helium  and  radium 
emanation.  The  epoch-making  hypothesis  which  fairly  met 
and  explained  every  known  fact  and  withstood  the  most 
searching  criticism  was  one  of  the  boldest  and  at  the  same  time 
the  simplest  and  most  comprehensive  ever  introduced  into  our 
science.  This  was  the  distintegration  hypothesis  of  Rutherford 
and  Soddy  (481).  These  scientists  were  at  the  time  professors 
in  the  department  of  physics  of  McGill  University,  Montreal. 

The  sketch  of  this  remarkable  hypothesis  already  given  (481) 
should  now  be  read  again.  According  to  this  hypothesis  radio- 
active atoms  are  more  or  less  unstable  systems.  The  first  step 
in  the  disintegration  of  a  radium  atom  is  the  throwing  off  of  an 


Radioactivity  and  the  Nature  of  Matter  579 

atom  of  helium  (an  alpha  particle)  and  the  leaving  behind 
of  a  residual  atom  of  smaller  mass,  an  atom  of  radium  emana- 
tion (Em), 

Ra-»He+Em  . 

If  we  suppose  that  the  helium  atom  shot  out  was  in  rapid  orbital 
motion  while  it  formed  a  part  of  the  radium  atom  we  have  at 
once  a  simple  explanation  of  the  high  velocity  of  the  alpha 
particle  and  therefore  of  the  energy  given  out  by  radium  as 
heat  (862),  since  the  kinetic  energy  of  the  alpha  rays  must  be 
changed  into  heat  when  the  rays  are  stopped  (368-371). 

866.  The  Radium  Series.  —  A  solid  body  of  any  kind  exposed 
to  radium  emanation  for  some  time  becomes  itself  strongly  radio- 
active, giving  out  alpha,  beta,  and  gamma  rays.  This  so-called 
excited  activity  decays  rather  rapidly  (to  half  in  about  thirty 
minutes).  Elaborate  experiments  have  shown  that  the  explana- 
tion of  these  facts  is  as  follows:  Just  as  a  radium  atom  disin- 
tegrates, throwing  off  an  alpha  ray  (helium  atom)  and  leaving 
an  atom  of  emanation,  so  an  atom  of  the  latter  in  turn  shoots  out 
another  helium  atom  and  leaves  a  new  and  lighter  residual  atom 
of  a  substance  called  radium  A,  having  a  very  short  period  (864). 
Radium  A  also  gives  out  alpha  rays  and  forms  radium  B,  and  the 
latter  in  turn  passes  into  radium  C.  The  series  of  radioactive 
products  of  radium  is  therefore  as  follows: 


The  excited  activity  obtained  from  the  emanation  consists 
of  the  products  A,  B,  and  C.  These,  together  with  the  emana- 
tion, are  present  in  radium  sealed  so  as  to  retain  the  emanation. 
Radium  itself  gives  only  alpha  rays.  The  beta  and  gamma  rays 
come  largely  from  radium  C. 

867.  The  Origin  of  Radium.—  Soon  after  the  discovery  of 
the  nature  of  radioactive  change  attempts  were  made  to  discover 
the  rate  of  decay  of  radium  itself.  Several  methods  were 
devised.  The  simplest  consisted  in  counting  the  number  of 
alpha  particles  given  off  by  a  known  amount  of  radium  in  a 
measured  period  of  time.  It  was  arranged  so  that  each  alpha 
particle  would  strike  a  zinc  sulfide  screen  and  produce  a  flash  of 


580  Introduction  to  General  Chemistry 

light  (859),  the  number  of  which  flashes  could  be  counted. 
Since  we  know  the  atomic  weight  of  radium  (857)  and  the  actual 
number  of  atoms  in  one  gram  atomic  weight  of  any  element 
(6.o6Xio23)  (194),  such  an  experiment  furnished  a  means  of 
calculating  the  rate  of  decay  of  radium.  This  rate  is  about 
0.04  per  cent  per  year,  which  corresponds  to  a  period  of  1,850 
years.  But  even  this  period,  long  as  it  is  in  terms  of  the  span 
of  human  life,  is  very  short  compared  with  the  age  of  the  earth 
as  judged  by  the  conclusions  of  geologists;  so  that  if  the  earth 
were  at  the  beginning  made  of  pure  radium,  after  even  a  million 
years  far  less  radium  would  remain  than  we  now  find  in  many 
radium-bearing  minerals.  But  as  the  age  of  the  earth  is  to  be 
reckoned  in  hundreds  of  millions  of  years,  and  as  there  is  a 
practical  certainty  that  the  proportion  of  radium  never  was 
high,  we  are  forced  to  the  conclusion  that  the  radium  now  found 
in  the  earth  is  being  formed  as  fast  as  it  decays.  When  these  con- 
clusions had  been  reached  it  became  a  matter  of  much  scientific 
interest  to  discover  the  origin  of  radium. 

The  parent  of  radium,  as  we  may  call  the  hypothetical  sub- 
stance from  which  we  may  suppose  radium  to  be  formed,  would 
very  likely  conform  to  the  following  specifications:  It  would  be 
associated  with  radium  in  minerals;  the  ratio  of  radium  to  the 
substance  would  be  constant  (for  the  same  sort  of  reason  that 
sealed  radium  contains,  after  a  sufficient  interval,  a  constant 
proportion  of  emanation) ;  the  parent  would  probably  be  radio- 
active, with  a  very  slow  rate  of  decay  compared  with  that  of 
radium;  the  atomic  weight  of  the  parent  would  exceed  that  of 
radium.  Only  one  element  conforms  to  all  these  specifications, 
namely  uranium  (836).  This  element  is  beyond  doubt  the  parent 
of  radium.  The  latter  element  is  found  only  in  uranium  minerals 
and  in  the  fixed  ratio  of  one  part  of  radium  to  three  million  parts 
by  weight  of  uranium. 

Uranium-bearing  minerals  always  contain  mechanically 
imprisoned  helium  which  has  been  formed  by  radioactive 
change.  In  fact  it  was  in  such  a  mineral,  cleveite,  that  helium 
was  discovered  (794). 


Radioactivity  and  the  Nature  of  Matter 


581 


868.  The  Uranium  Series. — The  change  of  uranium  into 
radium  is  not  direct  but  takes  place  by  the  intermediate  forma- 
tion of  a  series  of  radioactive  products,  as  indicated  in  Table  LIV, 
which  also  includes  the  entire  radium  series. 

TABLE  LIV 
THE  URANIUM-RADIUM  SERIES  (SLIGHTLY  CONDENSED) 


Element 

Period 

Rays 

Atomic  Weight 

Uranium  
Uranium  Xi  
Uranium  X2  
Uranium  2  
Ionium  
Radium  
Emanation  
Radium  A 

5X10'  years 
24  .  6  days 
69  sec. 
2Xio6  years 
2X108  years 
1,850  years 
3  .  85  days 
3  o  min. 

a 
0 

a 
a 
a 
a 

238 
234 
234 
234 
230 
226 
222 

218 

B 

26  8  min. 

0 

214 

c 

IQ  <  min. 

a,  /3    7 

214 

D  
E  
F  
G     .    .  .  . 

12  years 
5  days 
136  days 

"ft" 
a 

2IO 
210 
210 
206 

869.  The  Atomic  Weights  of  Radioactive  Substances. — If 

the  disintegration  hypothesis  is  correct  the  atomic  weight  of  a 
product  formed  by  an  alpha-ray  change  should  be  less  than  that 
of  its  parent  by  the  atomic  weight  of  helium,  namely  4  units. 
By  reason  of  the  very  small  mass  of  an  electron  (479)  the  atomic 
weight  is  practically  unchanged  as  the  result  of  a  beta-ray 
change.  Taking  the  atomic  weight  of  uranium  (836)  in  round 
numbers  as  238,  the  calculated  atomic  weights  of  its  products 
are  those  given  in  the  last  column  of  Table  LIV  (868).  We  can 
easily  test  these  conclusions  in  one  case,  that  of  radium.  The 
calculated  value  is  226.2,  if  U=  232.2,  while  Mme  Curie  found  by 
experiment  the  nearly  identical  value  226.  Therefore  it  appears 
probable  that  the  calculated  values  given  in  the  table  are 
essentially  correct. 

870.  Thoriu m  and  Its  Active  Products.— The  element  thorium 
(820,  834)  is  radioactive  with  an  intensity  of  activity  about  equal 
to  that  of  uranium.    Like  the  latter,  the  former  gives  rise  to  a 


582  Introduction  to  General  Chemistry 

long  series  of  active  products,  of  which  the  first,  mesothorium, 
is  the  only  one  of  technical  importance.  Mesothorium  is  identical 
with  radium  in  its  chemical  behavior,  and  no  means  is  known  of 
separating  a  mixture  of  the  two.  They  are  readily  distinguished 
by  their  periods  and  by  their  radioactive  products.  Meso- 
thorium has  a  period  of  5.5  years  as  against  1,850  years  for 
radium.  Mesothorium  itself  gives  no  rays  but  changes  into  a 
series  of  active  products  giving  alpha,  beta,  and  gamma  rays. 
Each  product  has  its  own  characteristic  chemical  properties, 
rays,  and  period.  One  of  these  products  is  radiothorium,  having 
a  period  of  two  years.  This  gives  out  alpha  rays  and  forms 
thorium  X,  which  in  turn  forms  a  gaseous  emanation  having  a 
period  of  less  than  one  minute.  Mesothorium  is  produced  as 
a  by-product  of  the  manufacture  of  thorium  and  is  used  as  a 
substitute  for  radium. 

871.  Isotopes. — A  most  unlooked-for  set  of  facts  was  brought 
to  light  by  the  investigation  of  the  chemical  behavior  of  certain 
radioactive  products.  We  have  mentioned  that  mesothorium 
and  radium  are  identical  in  chemical  behavior  (870).  Thorium 
X  is  also  identical  chemically  with  these  two,  and  no  means  is 
known  of  separating  mixtures  of  this  substance  and  either 
radium  or  mesothorium.  Substances  identical  in  chemical 
behavior  are  called  isotopes. 

Several  groups  of  isotopes  are  known.  Thus  radiothorium 
and  uranium  X  are  iso topic  with  thorium,  since  in  every  chemical 
reaction  they  both  behave  qualitatively  and  quantitatively 
exactly  like  thorium.  Each  of  the  three  can  be  obtained 
separately,  but  if  we  mix  the  separate  substances  it  is  impossible, 
by  any  known  means,  to  separate  them.  The  final  product  of 
the  radium  series,  formed  by  the  decay  of  RaF  and  known  as 
RaG  (at.  wt.  206),  is  an  inactive  substance  iso  topic  with  lead 
(at.  wt.  207.2),  and  curiously  enough  RaD  (at.  wt.  210)  is  also 
iso  topic  with  lead  and  therefore  also  with  RaG. 

Isotopes  are  not  identical  in  all  their  properties;  otherwise 
they  would  be  identical  substances  and  consequently  indis- 
tinguishable. Thus,  for  example,  radium  differs  from  meso- 
thorium in  the  following  respects:  Radium  has  a  period  of 


Radioactivity  and  the  Nature  of  Matter 


583 


1,850  years;  mesothorium  has  a  period  of  5 .5  years.  Ra  gives 
alpha  rays;  Ms  gives  no  rays.  Ra  changes  into  a  gas,  Em, 
which  gives  alpha  rays  and  has  a  period  of  3.85  days;  Ms 
changes  into  a  solid  which  gives  beta  rays  and  has  a  period  of 
6.2  hours.  Ra  has  an  atomic  weight  of  226.0,  while  Ms  has 
the  value  228.4.  No  difference  in  the  spectra  of  isotopes  has 
yet  been  discovered. 

872.  The  Valence  of  Radioactive  Substances. — Although  but 
very  few  of  the  radioactive  products  of  the  uranium-radium 
series  (Table  LIV)  or  the  thorium  series  have  as  yet  been  obtained 
in  weighable  amounts  in  pure  form,  still  we  know  much  about 
the  chemical  and  physical  properties  of  these  substances  and 
are  thus  able  to  learn  their  valence.  For  example,  when  we 
know  that  mesothorium  is  isotopic  with  radium  and  that  ths 
valence  of  the  latter  is  two,  we  are  safe  in  concluding  that 
mesothorium  has  the  same  valence.  Since  UXZ  and  ionium  are 
isotopic  with  thorium  we  conclude  that  all  three  have  the  same 
valence,  namely  four;  and  since  U2  is  isotopic  with  uranium  itself 
its  valence  is  six  (836).  These  examples  suffice  to  illustrate 
the  principle  employed.  Let  us  now  consider  the  first  part  of 
the  uranium-radium  series.  In  Table  LV  the  first  line  gives  the 

TABLE  LV 


U 

ux« 

ux, 

Ua 

lo 

Ra 

Em 

Rays     

a 

j8 

0 

a 

a 

a 

a 

Valence  

6 

4 

<; 

6 

4 

2 

o 

» 

symbols  of  the  products  of  uranium  in  the  order  of  their  forma- 
tion (868);  the  next  line  shows  the  kind  of  rays  given  out  by 
each  substance;  while  the  third  line  shows  the  valence  of  each. 
Fajans  discovered  a  remarkable  relation  between  the  facts 
set  forth  in  this  table,  namely,  (i)  the  valence  of  the  product  of 
an  alpha-ray  change  (858)  is  less  by  two  than  that  of  the  parent; 
(2)  the  valence  of  the  product  of  a  beta-ray  change  (860)  is  greater 
by  one  than  that  of  the  parent.  This  statement  is  called  Fajans' 
Law.  With  a  little  amplification  the  law  applies  to  all  radio- 
active substances  and  their  transformations. 


584  Introduction  to  General  Chemistry 

873.  The  Meaning  of  the  Term  Element. — A  direct  defini- 
tion of  the  term  element  has  not  been  given  in  this  text  (see  31) . 
To  say  that  an  element  is  a  substance  incapable  of  further  decom- 
position would  exclude  uranium,  radium,  and  thorium  as  well 
as,  of  course,  all  other  radioactive  substances  from  the  list  of 
elements;  and  this  would  certainly  be  an  unwise  classification. 
Since  radioactive  changes  are  always  spontaneous  and  entirely 
beyond  human  control,  it  would  seem  best  to  define  the  term 
element  thus:  An  element  is  a  substance  which  cannot  be  decom- 
posed into  simpler  substances  at  the  will  of  the  experimenter.    We 
shall  then  class  all  radioactive  products  (even  those  like  thorium 
emanation,  period  53  seconds)  as  elements.     All  these  "new" 
elements  have  been  satisfactorily  fitted  into  enlarged  periodic 
tables.     For  example,  see  the  lower  part  of  the  Harkins  table 
(849).     It  will  be  seen  that  several  elements  occupy  the  same 
position.    Except  for  the  rare  earths  these  are  isotopes. 

874.  The  Technical  Production  of  Radium. — Radium  was 
first  made  in  considerable  quantities  from  pitchblende,  a  mineral 
having  the  composition  U3C>8.     This  mineral  is  so  rare  that  it  is 
no  longer  a  satisfactory  source  of  radium.    The  bulk  of  the 
radium   now  produced   is   made    from   carnotite,    a    mineral 
of  somewhat  variable  composition,   but  containing  uranium, 
vanadium,  and  potassium  as  its  usual  constituents.     This  ore 
is  found  chiefly  in  southwestern  Colorado  and  adjoining  regions 
of  Utah,  for  the  most  part  as  cementing  material  in  sandstone. 
The  ordinary  ore  as  mined  carries  less  than  i  per  cent  of  uranium 
and  therefore  but  one  part  of  radium  in  300  million  parts  by 
weight  of  ore!    There  are  several  ways  of  extracting  the  radium, 
but  all  of  them  are  based  on  the  fact  that  the  radium  present 
remains  with  the  barium  present  when  the  latter  is  separated. 
If  insufficient  barium  is  present  in  the  ore,  enough  barium 
chloride  is  added  to  bring  the  proportion  up  to  about  i  per  cent. 
Usually  the  barium  (with  the  radium)  is  first  obtained  as  sulf  ate 
(164).    This  salt  is  converted  into  chloride  and   the  latter 
subjected  to  fractional  crystallization  in  the  manner  already 
described  (856).    A  much  more  rapid  separation  of  barium  and 
radium  results  from  the  crystallization  of  the  bromides,  since 


Radioactivity  and  the  Nature  of  Matter  585 

in  this  case  above  95  per  cent  of  the  radium  separates  with  half 
of  the  barium  salt.  In  another  process  the  hydroxides  of  barium 
and  radium  are  crystallized,  in  which  case  the  radium  concen- 
trates in  the  mother-liquors  instead  of  in  the  crystals. 

The  world's  total  production  of  radium  up  to  the  end  of  1918 
was  between  100  and  125  grams,  over  half  of  which  was  made 
in  the  United  States  of  America. 

875.  How  Radium  Is  Measured  and  Sold. — Radium  usually 
comes  on  the  market  as  an  isomorphous  (810)  mixture  of  its 
bromide  with  about  an  equal  amount  of  barium  bromide.     The 
selling  price  of  radium  is  always  based  on  the  radium  content  of 
material,  stated  in  milligrams  of  radium  element.     The  radium 
content  of  a  given  sample  is  not  found  by  the  use  of  a  balance 
but  is  determined  by  the  ionizing  power  of  the  gamma  rays  (86 1) 
as  shown  by  an  electroscope  (770,  854).     Comparative  measure- 
ments are  always  made  with  a  tube  of  radium  salt  of  exactly 
known   radium    content    (a   standard).     Every   purchaser   of 
radium,  before  paying  for  it,  can  have  the  material  measured  by 
the  United  States  Bureau  of  Standards  and  get  a  certificate  for 
a  small  fee.    During  the  year   1918   the  price  was  close  to 
$120.00  per  milligram  of  radium  element.     This  is  more  than 
100,000  times  the  value  of  gold!     But  when  it  is  known  that 
a  ton  of  ordinary  ore  yields  only  two  milligrams  of  radium  and 
that  it  requires  three  to  four  months'  work  to  extract  and  refine 
the  radium,  the  high  price  is  readily  understood. 

876.  The  Use  of  Radium  in  Therapeutics.— The  term  thera- 
peutics is  defined  as  being  that  branch  of  medical  science  that 
treats  of  the  action  of  remedial  agents  on  the  human  body.     The 
beta  and  gamma  rays  of  radium  are  capable  of  producing  serious 
"burns"  of  the  skin  and  underlying  tissues.     The  effect  produced 
is  proportional  to  the  quantity  of  radium  and  the  length  of  time 
of  the  exposure.     Healthy  tissue  is  much  less  affected  by  these 
rays  than  is  abnormal  and  unhealthy  tissue.     It  is  by  reason  of 
this  fact  that  radium  finds  an  important  use  as  a  therapeutic 
agent.    Exposure  of  certain  kinds  of  abnormal  growths  on  the 
body  to  radium  rays  results  in  their  destruction  and  subsequent 
removal.     Some  forms  of  cancer  respond  very  favorably  to  such 


586  Introduction  to  General  Chemistry 

treatment.  In  cases  where  large  masses  of  cancerous  tissue  are 
removed  by  surgical  operation  treatment  with  radium  rays  is 
often  subsequently  employed  to  destroy  the  remaining  portions 
not  accessible  to  the  knife.  In  hopeless  cases,  where  a  cure 
cannot  be  expected,  radium  treatment  is  useful  in  greatly 
alleviating  pain.  In  most  respects  radium  rays  produce  thera- 
peutic effects  like  those  due  to  X-rays.  The  great  advantage 
over  X-rays  in  the  use  of  radium  arises  from  the  fact  that  the 
radium  container  is  so  minute  that  it  can  be  applied  exactly 
where  its  rays  are  required  to  act,  a  procedure  impossible  with  an 
X-ray  bulb.  Radium  in  considerable  quantity  (100  mg.  or  more) 
may  be  safely  handled  if  contained  in  thick- walled  lead  tubes. 

877.  X-Ray  Spectra. — Ordinary  visible  light,  which  is  made 
up  of  ether  vibrations  of  various  wave-lengths  (788),  can  be 
spread  out  into  a  spectrum  of  its  component  colors  in  another 
way  besides  that  by  the  use  of  a  prism,  namely  by  means  of  a 
diffraction  grating.  The  latter  consists  of  a  polished  plane 
surface,  as  of  glass  or  metal,  ruled  with  an  enormous  number  of 
fine  parallel  lines,  often  as  little  as  one-thousandth  of  a  milli- 
meter apart.  The  principle  of  the  production  of  spectra  by 
such  gratings  is  discussed  in  most  textbooks  on  physics. 

For  many  years  after  the  discovery  of  X-rays  numerous 
attempts  were  made  to  obtain  X-ray  spectra,  but  all  without  any 
success,  by  the  use  of  either  prisms  or  gratings.  It  then  occurred 
to  Laue  that  if  X-rays  consisted  of  very  much  shorter  waves  than 
visible  light  it  would  be  impossible  to  rule  gratings  with  lines 
sufficiently  fine  and  close  together  to  show  the  expected  effect. 
In  fact,  he  calculated  that  if  X-rays  consist  of  waves  one- 
thousandth  as  long  as  those  of  visible  light,  in  order  to  get  their 
spectrum  it  would  require  a  grating  with  lines  no  farther  apart 
than  the  diameters  of  ordinary  atoms  (196).  To  rule  such  a 
grating  would  be  a  task  beyond  human  skill.  Then  came  the 
brilliant  idea  that  nature  makes  such  gratings  on  every  hand, 
since  in  every  crystal  the  molecules  are  arranged  in  rows  and 
layers  like  bricks  in  a  wall  (204).  Several  experimenters  applied 
Laue's  ideas  and  soon  brought  forth  a  great  number  of  new  and 
interesting  facts,  a  few  of  which  we  shall  discuss  briefly. 


Radioactivity  and  the  Nature  of  Matter 


587 


When  X-rays  from  a  given  source  strike  the  flat  surface  of 
a  crystal  obliquely  the  reflected  rays,  brought  to  focus  on  a 
photographic  plate,  produce  a  well-defined  line  spectrum  con- 
sisting of  a  small. number  of  lines.  Each  line  corresponds  to  a 
definite  wave-length  (just  as  ia  the  case  of  visible  light) ,  which 
length  can  be  easily  calculated. 

878.  Atomic  Numbers. — The  X-ray  spectrum  given  by  one 
element  used  as  the  target  of  the  X-ray  bulb  is  characteristic  for 
that  element  and  different  from  that  of  any  other  element.  In  this 
connection  a  remarkably  simple  relation  was  discovered  in  1914 
by  the  young  English  physicist  Moseley:  the  wave-length  of  the 

TABLE  LVI 
PERIODIC  TABLE  ACCORDING  TO  ATOMIC  NUMBERS,  H:i 


Pe- 
riod 

0 

I 
A          B 

II 
A          B 

III 
A         B 

IV 
A          B 

V 
A          B 

VI 
A          B 

VII 
A          B 

VIII 

I 

2 

He 

Li3 

4 
Be 

5 
B 

6 

c 

N 

8 

O 

2 

10 

Ne 

IO 

Na 

12 

Mg 

li 

14 
Si 

V5 

16 

g 

17 

a 

18 
A 

K19 

20 
Ca 

21 

Sc 

22 

Ti 

23 
V 

24 
Cr 

Mn5 

26  27  28 
Fe  Co  Ni 

3 

29Cu 

3°Zn 

3IGa 

32Ge 

33 

As 

34  Se 

35 
Br 

36 
Kr 

Kb" 

38 
Sr 

Y39 

40 

Zr 

cb3 

42 

Mo 

?43 

44  45  46 
Ru  Rh  Pd 

4 

4?Ag 

48 
Cd 

49 
In 

5° 
Sn 

51  Sb 

52 
Te 

53 

54 
Xe 

55 
Cs 

56 
Ba 

57 
La 

& 

B 

60 
Nd 

61 
? 

5 

62 
Sa 

F3 

Eu 

64 
Gd 

& 

66 
Dy 

67 
Ho 

67 
Er 

69 
Tm 

70 

Yb 

72 
Lu 

73 
Ta 

w74 

75 

76  77  78 
Os  Ir  Pt 

79 
Au 

80 
Hg 

81 
Tl 

82 
Pb 

83  Bi 

84 

85 

7 

86 

Nt 

?  8? 

—  •  — 

88 
Ra 

89 
•> 

90 
Th 

91 

u92 

588  Introduction  to  General  Chemistry 

most  intense  line  decreases  quite  uniformly  as  we  pass  from  one 
element  to  the  next  higher  in  atomic  weight.  Moseley  also 
showed  that  if  the  elements  are  assigned  numbers,  called 
atomic  numbers,  in  the  order  of  their  increasing  atomic  weights 
there  is  a  very  simple  numerical  relation  between  the  atomic 
number  of  an  element  and  the  wave-length  of  the  strongest  line 
of  its  X-ray  spectrum.  The  atomic  numbers  calculated  from 
the  wave-lengths  as  found  by  experiment  differed  from  whole 
numbers  by  not  more  than  i  or  2  per  cent,  and  this  difference 
could  easily  be  ascribed  to  error  of  experiment.  The  only  cases 
where  the  atomic  numbers  so  found  did  not  follow  exactly  the 
order  of  increasing  atomic  weights  were  met  with  in  those  pairs 
of  elements  (A  and  K,  Te  and  I)  where  a  departure  from  the 
usual  order  is  necessary  to  bring  the  elements  into  their  proper 
groups  in  the  Periodic  Table  (829).  In  Table  LVI  the  elements 
are  arranged  in  a  periodic  table  strictly  in  the  order  of  their 
atomic  numbers.  The  anomalies  (848)  now  disappear  com- 
pletely. It  would  therefore  appear  that  the  atomic  number  of 
an  element  is  its  most  fundamental  constant.  In  Table  LVI 
the  rare  earth  elements  (820, 846)  are  included  within  the  heavy 
lines.  These  elements  form  a  single  family  of  Group  III  (849) 
and  do  not  belong  to  the  other  groups  indicated  by  their 
positions  in  this  table. 

879.  The  Structure  of  Crystals.— Professor  W.  H.  Bragg  and 
his  son  W.  L.  Bragg,  English  physicists,  have  followed  Laue's 
ideas  in  another  direction  and  obtained  a  great  deal  of  informa-. 
tion  regarding  the  arrangement  of  molecules  and  atoms  in  a 
great  variety  of  crystals.     The  simplest  way  of  investigating  the 
structure  of  a  crystal  consists  in  getting  the  impression  made  on 
a  photographic  plate  when  a  small,  round  beam  of  X-rays  passes 
through  a  thin  slice  of  the  crystal  and  strikes  the  plate  some 
inches  beyond.     The  reproduction  of  such  a  photograph  is  shown 
in  Fig.  121.     From  the  nature  of  the  pattern  produced  the  geo- 
metrical arrangement  and  distances  from  one  another  of  the 
atoms  of  a  crystal  can  be  calculated. 

880.  Atomic  Structure  and  the   Nature  of  Matter.— The 
modern  view  of  the  structure  of  an  atom  has  been  described 


Radioactivity  and  the  Nature  of  Matter  589 

briefly  in  section  470.    The  facts  epitomized  in  the  periodic  law, 
together  with  the  phenomena  of  radioactivity,  show  very  con- 
vincingly that  atoms  must  be  constructed  according  to  very 
definite  plans,  and  that  the  arrangement  of  the  parts  (electrons, 
atoms  of   helium,   and    the 
atomic  nucleus,  470)  in  the 
atom  of  an  element  must  de- 
termine its  properties.    Some 
of  the  world's  ablest  physi- 
cists and  chemists  have  de- 
voted much  attention  to  the 
problems  of  atomic  structure, 
and  as  a  result  several  more 
or  less  definite  hypotheses 
have  resulted,  each  aiming  to 
explain  and  correlate  as  many 
facts  as  possible.    As  yet  none 
of  these  suggestions  is  wholly 
satisfactory,  although  several 

of  them  account  very  well  for  many  of  the  facts  with  which 
we  are  familiar. 

On  one  point  there  is  general  agreement:  namely,  that  the 
atomic  number  of  an  element  represents  the  number  of  electrons 
encircling  the  electro-positive  nucleus  of  the  atom.  Since  atoms,  as 
such,  are  electrically  neutral,  the  positive  charge  of  the  nucleus 
must  also  be  proportional  to  the  atomic  member.  It  seems  prob- 
able that  the  nucleus  also  contains  some  electrons  and  that  its 
positive  charge  represents  the  excess  of  its  positive  over  its  nega- 
tive electricity. 

In  changes  of  atoms  into  ions  and  vice  versa  (oxidations  and 
reductions,  501-507)  the  valence  electrons  lost  or  gained  are  those 
of  the  outer  ring,  and  not  of  the  nucleus.  In  radioactive  changes 
the  alpha  and  beta  rays  doubtless  come  from  the  nucleus,  so  that 
the  new  residual  atom  has  a  nucleus  different  in  positive  charge 
from  its  parent  and  therefore  surrounded  by  a  different  number 
of  electrons  in  the  outer  rings.  It  is  on  this  basis  that  Fajans' 
Law  (872)  receives  its  explanation. 


590  Introduction  to  General  Chemistry 

It  is  of  interest  to  note  that  the  chemist's  control  over  the 
composition  of  the  atom  is  limited  to  the  removal  and  replace- 
ment of  some  of  the  electrons  of  the  outer  rings;  changes  of  the 
nucleus  are  solely  the  result  of  spontaneous  radioactive  pro- 
cesses entirely  beyond  human  control. 

Since  the  positive  charge  of  the  nucleus  is  the  excess  of  its 
positive  over  its  negative  electricity,  it  is  possible  for  two  or  more 
differently  composed  nuclei  to  have  equal  nuclear  charges. 
Atoms  containing  nuclei  differing  thus  would  have  equal 
numbers  of  electrons  in  the  outer  rings  and  therefore  be  identi- 
cal in  chemical  but  not  in  radioactive  properties.  In  conse- 
quence it  would  be  possible  to  have  two  (or  more)  different 
elements  with  identical  chemical  properties,  in  other  words 
isotopes  (871). 

In  conclusion,  it  may  be  stated  that  all  chemists  and 
physicists  are  now  agreed  that  the  elements  are  but  modifications 
of  the  same  primitive  forms  of  material,  including  in  the  latter 
term  electrons  and  whatever  else  may  compose  an  atom.  It 
may  well  be  that  all  matter  is  entirely  of  an  electrical  nature. 
This  view  has  led  some  persons  to  declare  that  "  there  is  no 
such  thing  as  matter,"  that  "  every  thing  is  electricity."  This 
statement  is  scarcely  warranted.  It  is  just  as  if,  when  we  have 
discovered  that  a  potato  is  composed  of  starch,  fiber,  and  water, 
we  should  declare  "there  is  no  such  thing  as  a  potato"!  As  for 
matter,  it  remains  what  it  always  has  been,  only  we  know  more 
about  it. 


CHAPTER  XXXIII 

METALLURGY 

881.  Introduction. — Metallurgy,  or  the  art  of  extracting 
metals  from  their  ores,  had  its  beginning  in  prehistoric  times. 
In  our  present-day  civilization  it  is  the  basis,  of  one  of  our 
greatest  industries.  The  colossal  importance  of  it  may  be 
judged  from  Table  LVII,  which  gives  the  annual  production 
of  metals  in  the  United  States,  the  values  of  the  output,  and 
the  percentage  of  the  world's  supply  furnished  by  the  United 
States.  In  the  cases  of  iron,  copper,  silver,  zinc,  lead,  and 
aluminum,  the  United  States  produces  more  than  any  other 
one  country.  The  total  value  of  its  annual  production  is  over 
two  billion  dollars. 

As  will  be  shown  later,  this  enormous  industry  may  well  be 
called  a  chemical  industry  since,  in  the  first  place,  the  science 
of  chemistry  has  contributed  greatly  to  improvements  practised 
in  the  general  procedures,  and  since,  in  the  second  place, 
virtually  all  of  the  enormous  output  of  metals  is  prepared  under 
supervision  from  chemical  laboratories. 

It  will  be  the  purpose  of  the  present  chapter  to  present  a 
brief  survey  of  the  metallurgy  of  the  more  important  metals, 
namely,  iron,  copper,  zinc,  lead,  and  aluminum,  and  incidentally 
certain  aspects  of  the  metallurgy  of  silver  and  gold. 

TABLE  LVII 
ANNUAL  PRODUCTION  OF  METALS  IN  THE  UNITED  STATES 


Year 

Amount 

Value 

Per  cent.  > 

Pig  iron  

1918 

42.000.000  tons 

$1,300,000,000              55 

Copper  

1917                944,000  tons 

500,000,000  1          60 

Zinc  

1917                685.000  tons 

120,000,000  i          32 

Lead  

1917 

580,000  tons 

100,000,000 

34 

Aluminum  

1918 

112,000  tons 

74,000,000 

46 

Gold  

1917 

4,051,000  oz. 

83,750,000 

20 

Silver  

1917 

71,740,000  oz. 

58,000,000 

42 

1  U.  S.  A.  production  as  percentage  of  total  for  the  world. 

591 


592  Introduction  to  General  Chemistry 

882.  Economic  Importance  of  Iron. — Of  all  metals,  iron  is 
the  most  important.     Not  only  is  this  true  because  -it  is  abun- 
dant, but  because  of  the  cheapness  of  its  production  and  its 
manifold  useful  properties.     To  appreciate  this  we  have  only 
to  reflect  that  most  of  our  earthly  possessions  which  do  not 
consist  wholly  or  in  part  of  iron  have  been  made  with  the  aid 
of  iron  machinery,  which  is  driven  by  engines  or  motors  con- 
sisting wholly  or  in  part  of  iron.     Iron  is  the  only  element  which 
is   decidedly   magnetic    (of   high   magnetic   permeability)  and 
so   is    the    essential   material   for   magnets.     Thus    dynamos, 
motors  and  other  electrical  machinery  which  are  the  basis  of 
the  modern  generation  and  use  of  electricity  would  not  be 
possible  without  iron.     It  is  easy  to  see  that  the  existing  material 
aspect  of  our  present-day  civilization  is  dependent  upon  this 
element.     It  is  therefore  not  strange  that  iron  is  produced  in 
far  greater  quantity  than  any  other  metal,  and  it  is  interesting 
to  find  that  the  tonnage  of  iron  produced  in  the  United  States 
is  about  twenty  times  that  of  all  the  other  metals  combined. 
(See  Table  LV1I.) 

883.  Commercial  Grades  of  Iron.    Steel. — Broadly  speak- 
ing, commercial  grades  of  iron  belong  to  one  of  three  classes: 
wrought  iron,  cast  iron,  and  steel.     All  grades  contain  small 
amounts    of    certain   other    elements    which   have   important 
bearing  on  the  properties  of  the  product.     Of  these  elements 
the  most  important  is  carbon.     Others  are  silicon,  manganese, 
phosphorus  and  sulfur. 

The  terms  wrought  iron  and  cast  iron  designate  substances 
of  more  definite  composition  than  does  the  term  steel,  which 
applies  to  products  of  certain  processes.  Some  kinds  of  steel 
are  nearly  identical  in  composition  and  properties  with  wrought 
iron.  The  latter  and  also  low  carbon  or  mild  steel  contain  less 
than  0.3  per  cent  of  carbon.  Half-hard  and  high  carbon  steel 
contains  from  0.3  per  cent  to  2.2  per  cent  of  carbon.  Cast 
iron  contains  from  2.2.  to  5  per  cent  of  carbon. 

Wrought  iron  and  mild  steel  are  moderately  soft  metals. 
High  carbon  steel  may  be  either  extremely  hard  or  rather  soft, 


Metallurgy  593 

according  to  the  heat  treatment  given  it.     Cast  iron  is  moder- 
ately hard  and  brittle,  but  can  be  machined  to  shape. 

884.  Iron  Ores.  —  It  seems  probable  that  iron  constitutes 
about  5  per  cent  of  the  earth's  crust  (F.  W.  Clarke).     The  only 
iron-bearing  minerals  of  importance  are  those  which  contain 
40  per  cent  or  more  of  this  element  and  only  small  percentages 
of  objectionable  impurities,  especially  phosphorus  and  sulfur. 
In   the  United  States  the  most  important  ore  is  hematite, 
Fe203,    (328).    Limonite,   a    hydrated    hematite    called    also 
brown  ore,  and  magnetite,  FesCU  (173)  are  also  mined  in  con- 
siderable amounts;  but  altogether  they  constitute  only  6  to  8 
per  cent  of  the  total  iron  ore  of  this  country. 

No  less  than  85  per  cent  of  all  the  iron  ore  mined  in  the 
United  States  comes  from  a  single  region  which  lies  to  the 
south  and  west  of  Lake  Superior,  in  adjacent  parts  of  the  states 
of  Minnesota,  Michigan  and  Wisconsin.  The  ore  of  the  Lake 
Superior  region  is  hematite.  The  next  most  important  dis- 
trict, located  in  Alabama,  produces  8  per  cent  of  the  total  ore. 
New  York  state  ranks  third  in  importance,  with  3  per  cent. 
The  remaining  states  produce  but  4  per  cent  of  the  total. 

885.  The  Chemistry  of  the  Metallurgy  of  Iron.—  The  princi- 
pal reactions  involved  in  the  conversion  of  ferric  oxide, 
(hematite),  into  metallic  iron  (329)  are 

Fe2O3+3C-»2Fe+3CO 
and 


Taken  together,  these  may  be  written: 

2Fe203+3C-*4Fe+3C02  (328) 

If  iron  ores  consisted  of  pure  oxides,  it  would  suffice  to  heat 
these  with  carbon  and  nothing  else.  However  in  practice  this 
is  not  the  case.  The  average  composition  of  Lake  Superior 
ore  for  the  year  1909  was  that  shown  in  Table  LVIII. 

It  thus  appears  that  components  other  than  iron  oxide  are 
present  in  large  amount  and  must  be  removed  in  the  metal- 
lurgical process.  The  elimination  of  silica,  alumina,  calcium, 
etc.  is  accomplished  by  converting  them  into  a  fusible  glass- 


594  Introduction  to  General  Chemistry 

TABLE  LVIII 

AVERAGE  PERCENTAGE  COMPOSITION  OF  LAKE  SUPERIOR  IRON  ORE  (RIES) 

Iron 58 . 45  per  cent 

Silica 7.67 

Alumina 2.23 

Calcium,  magnesium  and  manganese i .  80 

Phosphorus 09 

Sulfur 06 

Loss  on  ignition 4.12 

like  product  called  slag,  which  at  a  white  heat  forms  a  fluid  that 
floats  on  molten  iron  like  oil  on  water  and  can  be  separated 
mechanically. 

In  the  formation  of  slag  we  have  a  typical  reaction  in  a  fused 
mass.  For  example,  calcium  oxide  is  a  base  forming  oxide. 
Silica  is  an  acid  forming  oxide.  Heated  together  they  form 
a  salt,  calcium  silicate,  just  as  acids  and  bases  do  in  water  solu- 
tion. The  salt  is  molten  at  the  temperature  of  the  furnace  and 
dissolves  many  other  oxides  and  salts.  The  following  is  rep- 
resentative of  the  percentage  composition  of  slag: 

Si02    A12O3     CaO     MgO     CaS     Remainder 
34          12         43  6         3  2 

Slag  is  therefore  a  silicate  of  aluminum,  calcium,  and  magnesium 
with  small  proportions  of  other  substances.  Since  but  few 
ores  contain  the  slag  forming  elements  in  ideal  proportions, 
it  is  customary  to  add  to  the  ore  sufficient  materials  from  other 
sources  to  make  up  the  deficiency.  Usually  it  is  only  necessary 
to  add  limestone  (150). 

We  shall  now  describe  the  smelting  of  iron  ore  as  actually 
carried  out  in  the  blast  furnace. 

886.  The  Blast  Furnace. — A  blast  furnace  is  represented  in 
Fig.  122.  This  is  a  huge  structure,  60  to  90  feet  in  height, 
built  of  iron  and  lined  with  fire  brick.  This  furnace  gets  its 
name  from  the  fact  that  a  blast  of  air  is  supplied  near  the  base 
at  a  pressure  of  about  15  pounds  per  square  inch  above  atmos- 
pheric pressure.  The  air  before  entering  the  furnace  is  heated 
to  about  700°  (corresponding  to  a  red  heat)  in  a  way  later  to 
be  described.  The  charge  or  burden,  consisting  of  ore,  lime- 
stone and  coke  (328,  630)  is  introduced  at  the  top  of  the  furnace, 


Metallurgy  595 

which  is  closed  by  a  cone  or  "bell."  In  a  typical  case,  a 
charge  of  9  tons  of  ore,  5  tons  of  coke  and  3  tons  of  limestone 
is  added  every  quarter  of  an  hour,  day  and  night,  week  in  and 
week  out,  for  many  months  at  a  time.  As  the  charge  passes 
downward  it  soon  becomes  hot  enough  to  react  so  that  by  the 
time  it  has  reached  the  level  of  the  blast  pipes  (tuyeres)  it 
consists  largely  of  molten  iron  and  slag.  The  former  collects 
in  the  lowest  part  of  the  furnace  as  a  white  hot  fluid,  on  which 
floats  the  equally  hot  but  somewhat  more  viscous  liquid  slag. 
From  time  to  time  slag  is  tapped  from  the  furnace  through  an 
opening  made  for  the  purpose.  At  in- 
tervals of  a  few  hours  the  iron  is  also 
drawn  off  at  a  lower  level  opening  and 
run  into  troughs  made  in  a  bed  of  sand. 
The  smaller  troughs,  about  3  feet  long 
and  4  inches  wide,  branch  off  from  the 
main  channel  at  uniform  distances.  The 
solidified  product  is  called  pig  iron.  A 
large  furnace  charged  as  above  indicated 
will  produce  500  tons  (1,000,000  pounds) 
of  pig  iron  in  24  hours. 
887.  The  Blast  Furnace  Plant. — In 

order  that  the  reduction  of  ore  to  metal          TAP  EOLB 

.    .  FIG.  122 

may  be  complete,  it  is  necessary  to  use 

more  carbon  (in  the  form  of  coke)  than  that  indicated  by  the 
chemical  equations  (885).  As  a  result,  the  gases  issuing  from 
the  top  of  a  furnace  contain  a  large  proportion  of  carbon  mon- 
oxide. A  typical  product  will  contain  25  per  cent  of  carbon 
monoxide,  13  per  cent  of  carbon  dioxide,  3  per  cent  of  hydrogen, 
and  5  7  per  cent  of  nitrogen.  This  gas  is  saved  and  used  as  fuel 
to  heat  the  air  used  for  the  blast  and  also  to  produce  the  power 
to  run  the  air  blowers.  For  the  latter  purpose  it  is  either 
burned  under  the  boilers  of  steam  engines  or,  in  more  modern 
plants,  it  is  used  directly  as  the  fuel  for  huge  gas  engines.  The 
blast  furnace  requires  therefore  a  large  accessory  equipment,  the 
whole  forming  the  blast  furnace  plant.  The  more  important 
items  of  this  plant  will  now  be  briefly  described.  The  gases 


596  Introduction  to  General  Chemistry 

escaping  from  the  furnace  carry  a  great  deal  of  dust.  This  is 
first  separated  and  collected  in  a  dust  catcher  consisting  of  an 
immense  cylindrical  vessel  with  a  conical  base  in  which  the 
dust  settles  out  to  a  large  extent.  The  gases  may  be  further 
purified  in  washers. 

For  heating  the  blast,  so-called  stoves  are  used,  of  which  three 
or  four  are  required  for  each  blast  furnace.  A  stove  is  a  cylin- 
drical furnace  nearly  half  as  large  as  the  blast  furnace.  It 
consists  of  a  central  chamber  in  which  the  gas  burns  surrounded 
by  a  honey-comb  like  arrangement  (checker-work)  of  fire 
brick,  through  which  the  flame  and  hot  gases  pass  before  they 
escape  at  the  smoke-stack.  After  a  short  time  the  stove 
becomes  very  hot.  The  supply  of  gas  is  then  turned  off  and 
the  blast  of  (cold)  compressed  air  is  led  through  the  hot  iron 
stove  whereby  the  air  is  raised  to  a  temperature  of  about  700°. 
This  hot  air  is  used  to  blow  the  blast  furnace.  While  one 
stove  is  being  used  to  heat  the  blast,  others  are  being  heated 
by  burning  gas. 

888.  Pig  Iron  and  Slag. — Pig  iron  is  by  no  means  pure  iron. 
Its  composition  is  variable.     It  contains  about  94  per  cent  of 
iron,  3  to  4  per  cent  of  carbon,  i  to  2  per  cent  of  silicon,  i 
per  cent  of  manganese,  o.i  to  2  per  cent  of  phosphorus,  and 
less  than  o.i  per  cent,  of  sulfur.     Pig  iron  is  so  brittle  that  the 
pigs  are  easily  broken  into  smaller  pieces  by  means  of  a  sledge 
hammer.     Small  pieces  can  be  ground  to  a  powder  in  a  mortar. 
In  appearance,  pig  iron  is  metallic  gray,  and  shows  a  crystalline 
or  granular  facture. 

Pig  iron  is  an  intermediate  product  which  is  used  to  make 
purer  forms  of  iron,  about  78  per  cent  of  the  production  going 
into  steel,  19  per  cent  into  cast  iron,  and  2^  per  cent  into 
wrought  iron. 

Slag  is  largely  waste  material  which  is  used  for  road  building 
and  for  ballast  for  railroads.  A  good  quality  of  cement  is  made 
from  slag.  Mineral  wool,  used  for  fire  and  acid-proof  packing, 
is  made  by  blowing  steam  through  molten  slag. 

889.  Cast  Iron. — Cast  iron  is  the  form  of  the  metal  used  in 
making  stoves,  steam  radiators,  street  lamp  posts,  water  and  gas 


Metallurgy 


597 


CHARGING  FLOOR 


mains,  fire  hydrants,  etc.  where  strength  and  rigidity  but  'riot 
flexibility  nor  elasticity  are  required.  Iron  castings  are  made 
by  pouring  the  molten  metal  into  moulds  made  in  sand  from 
wooden  patterns  of  the  form  to  be  reproduced.  The  metal  is 
melted  in  a  cupola  furnace,  Fig.  123.  This  furnace  consists  of  a 
cylindrical  sheet  steel  shell  3  to  6  feet  in  diameter  and  1 5  to  2  5  feet 
high  lined  with  fire-brick.  The  bottom  is  covered  with  a  thick 
layer  of  sand.  A  tapping  hole  for  the  iron  is  located  at  the 
upper  level  of  the  sand  layer  and  on  the  opposite  side  of  the 
cupola  a  little  higher  up  is  another  tap  hole  for  slag.  Several 
inlets  (tuyeres)  for  the  air  blast 
are  provided.  Thus  the  cupola 
furnace  resembles  a  very  small 
blast  furnace.  Its  charge  consists 
of  pig  iron  and  coke  together  with 
enough  limestone  to  form  a  slag 
with  the  silica  of  the  coke  ash  and 
a  certain  additional  amount  for- 
med by  oxidation  of  a  small  part 
of  the  silicon  of  the  pig  iron  used. 
Usually  some  old  scrap  iron, 
broken  and  worn  out  castings, 
etc.  are  added  to  the  charge.  In 
the  furnace  the  coke  burns  in  the 
strong  air  blast  and  raises  the  temperature  above  the  melting 
point  of  the  pig  iron  (about  1100°).  The  molten  iron,  cov- 
ered by  a  layer  of  molten  slag,  collects  in  the  bottom  of  the 
furnace.  The  white  hot  liquid  iron  is  drawn  off  through  the 
tap  hole  into  a  ladle  from  which  it  is  poured  into  the  mould. 
An  iron  casting  or  moulding  establishment  is  called  a  foundry. 
There  are  many  grades  of  cast  iron  which  merge  impercep- 
tibly into  one  another.  The  differences  are  due  to  two  causes: 
the  chemical  composition  and  the  heat  treatment  the  product 
has  undergone.  In  addition  to  a  little  manganese  and  phos- 
phorus and  a  minute  amount  of  sulfur,  cast  iron  contains  be- 
tween %  and  2^  per  cent  of  silicon  and  2%  to  3^  Per  cent  °f 
carbon.  It  is  the  last  two  elements  in  particular  that  determine 


.TAP   HOL8 


FIG.  123 


598 


Introduction  to  General  Chemistry 


the  character  of  the  cast  iron.  The  two  principal  sorts  of  cast 
iron  are  designated  as  white  and  gray  respectively.  If  a  molten 
iron  low  in  a  silicon  (e.g.  %  per  cent)  and  containing  about  2% 
per  cent  of  carbon  is  cooled  rapidly,  white  cast  iron  results. 
This  is  silver  white  on  the  fracture  and  very  brittle.  It  is  so 
hard  that  it  cannot  be  machined  to  shape  in  a  lathe.  On  the 
other  hand,  iron  high  in  silicon  (e.g.  2>£  per  cent)  and  carbon 
(3^  Per  cent)  if  cooled  slowly  forms  gray  cast  iron,  gray  in 
color,  and  soft  enough  to  be  easily  machined.  The  reason  for 
these  facts  will  be  considered  in  a  later  section. 

890.  Wrought  Iron. — Wrought  iron  is  one  of  the  purest 
commercial  forms  of  the  metal.  It  contains  less  than  0.3  per 
cent  of  carbon  and  only  traces  of  silicon,  phosphorus  and  sulfur; 


FIG.  124 

although  it  holds  in  mechanical  admixture  a  small  amount  of 
solid  slag.  The  conversion  of  molten  pig  iron  into  wrought  iron 
is  accomplished  by  oxidizing  the  elements  to  be  eliminated, 
thus  converting  the  carbon  into  gaseous  monoxide,  CO,  and 
the  others  into  a  molten  slag  consisting  largely  of  ferrous  silicate 
and  ferrous  phosphate.  The  oxidation  is  carried  out  by  mixing 
the  molten  pig  iron  at  a  white  heat  with  oxides  of  iron. 

This  process  of  making  wrought  iron,  usually  known  as  pig 
boiling,  is  carried  out  in  a  so-called  puddling  furnace,  Fig.  1 24. 
This  is  a  small,  simple  form  of  reverberatory  furnace.  This 
kind  of  furnace  is  so  called  because  heat  from  the  flame  pro- 
duced in  the  fire  box  is  reflected  or  reverberated  from  the  top  of 
the  furnace  to  the  material  on  the  hearth  beneath.  In  a  rever- 
beratory furnace  the  material  to  be  heated  is  not  contaminated 
by  contact  with  the  fuel.  In  the  puddling  furnace  the  hearth 


Metallurgy  599 

is  lined  with  iron  oxides,  which  may  be  very  pure  ore  (hematite 
or  magnetite)  or  so-called  hammer  scale,  which  is  chiefly 
Fe?O4.  The  process  consists  in  melting  about  400  pounds  of 
pig  iron  on  the  hearth  and  adding  to  the  molten  metal  powdered 
iron  oxides  and  stirring  (rabbling)  the  mixture  thoroughly 
with  an  iron  bar.  A  vigorous  reaction  soon  begins.  The  car- 
bon monoxide  formed  by  the  reactions 

Fe,04+4C->4CO+3Fe 


escapes,  causing  the  material  to  appear  to  boil,  hence  the  term 
pig  boiling.  The  ferrous  oxide  formed  in  the  last  reaction 
unites  with  silicon  dioxide  formed  by  oxidation  of  the  silicon 
to  form  a  silicate  and  with  phosphorus  pentoxide  produced  by 
oxidation  of  the  phosphorus  to  form  a  phosphate.  These  fused 
salts  form  the  bulk  of  the  slag.  The  sulfur  present  is  partly 
converted  into  sulfide  that  passes  into  the  slag,  and  partly  into 
sulfur  dioxide,  which  escapes  as  a  gas. 

The  purified  iron  that  results  from  these  reactions  melts  at 
a  much  higher  temperature  than  does  pig  iron,  with  the  result 
that  as  the  process  approaches  completion  semi-solid  masses  of 
iron  form  in  the  mixture.  These  are  gathered  into  a  white  hot 
ball  by  the  puddler  and  worked  about  on  the  hearth  by  means 
of  an  iron  rod  (rabble)  until  the  mass  is  nearly  freed  from  slag. 
It  is  next  subdivided  and  the  smaller  plastic  lumps,  still  white 
hot,  removed  and  worked  by  hammers  or  mechanically  operated 
squeezers  to  remove  as  much  as  possible  of  the  still  liquid  slag. 
The  product  is  then  rolled  into  bars,  etc. 

Wrought  iron  is  exceedingly  tough  and  not  at  all  brittle. 
When  it  is  heated  to  a  high  temperature  it  is  easily  welded  (803), 
or  it  may  be  wrought  (forged)  into  any  desired  shape  on  a  black- 
smith's anvil. 

891.  The  Bessemer  Process.  —  The  manufacture  of  wrought 
iron  is  both  laborious  and  expensive.  This  is  because  it  re- 
quires most  strenuous  manual  labor  and  that  by  a  skilled  work- 
man to  produce  in  the  course  of  several  hours  a  batch  of  a  few 
hundred  pounds  of  this  desirable  form  of  iron.  Imagine,  then, 


6oo 


Introduction  to  General  Chemistry 


the  importance  of  a  process  that  produces  from  pig  iron  a  prod- 
uct equal  or  superior  to  wrought  iron  in  quality  in  batches  of 
15  to  20  tons,  with  little  hand  labor,  in  20  minutes'  time,  using 
the  cheapest  of  all  oxidizing  agents,  namely,  air! 

This  process  was  first  carried  out  by  an  American,  William 
Kelly,  in  1846,  and  later  perfected  after  many  years  of  per- 
sistent experimenting  by  Sir  Henry  Bessemer,  an  Englishman. 
Theoretically  the  process  is  very  simple:  a  powerful  blast  of  air 
is  forced  through  molten  pig  iron  contained  in  a  vessel  lined 
with  silica  or  silicates.  The  air  oxidizes  the  carbon,  silicon 
and  manganese  (889)  of  the  melted  iron  to  oxides.  The 
oxide  of  carbon,  being  a  gas,  passes  off,  the  silicon  dioxide 

unites  with  the  oxides  of  iron  and 
manganese  (as  in  the  pig  boiling 
process)  to  form  a  fusible  slag. 
The  iron  is  thus  freed  from  other 
elements  excepting  sulfur  and 
phosphorus.  By  a  modification 
of  the  process,  even  the  latter  can 
be  removed.  During  the  process 
much  heat  is  rapidly  produced 
and  the  temperature  of  the  ma- 
terials rises  so  high  that  both 
metal  and  slag  remain  molten. 
Simple  as  it  appears  when  thus 
stated,  enormous  difficulties  had 

to  be  overcome  before  the  process  was  perfected.  The  product 
is  known  as  Bessemer  steel. 

The  process  is  carried  out  in  a  huge  vessel  called  a  converter, 
Fig.  125.  This  is  built  of  steel  plates  and  lined  with  silicious 
material.  It  is  supported  on  trunions,  so  that  it  can  be  tilted 
to  pour  out  its  contents.  The  pig  iron  used  for  the  charge  is 
not  allowed  to  cool  after  being  drawn  from  the  blast  furnace 
(887)  but  is  swiftly  carried  in  a  great  ladle  mounted  on  a  rail- 
road truck  and  run  into  the  converter  and  blown  at  once. 
Silicon  and  manganese  burn  first;  after  a  few  minutes,  when 
the  oxidation  of  the  carbon  begins,  a  flame  shoots  out  to  a 


FIG.  125 


Metallurgy  60 1 

height  of  20  or  30  feet  with  a  great  shower  of  brilliant  sparks. 
After  15  or  20  minutes  the  flame  dies  down,  all  the  carbon 
having  been  burned  out.  At  this  stage  a  small  but  definite 
proportion  of  an  iron-manganese  alloy  called  ferro-manganese 
is  added  This,  by  reason  of  its  manganese  content,  removes 
excess  oxygen  or  reduces  oxides  of  iron  in  the  metal  and  thus 
prevents  brittleness  at  a  red  heat  (red  shortness)  that  would 
otherwise  exist.  The  carbon  content  of  the  product  is  brought 
up  to  the  desired  point  by  adding  anthracite  coal  or  other 
form  of  carbon,  and  the  metal  is  then  cast  in  iron  molds  (ingot 
molds)  about  15  inches  square  and  4.  or  5  feet  deep. 

The  whole  process,  from  the  time  the  molten  pig  iron  is 
drawn  from  the  blast  furnace  until  it  has  been  cast  into  ingots, 
requires  less  than  an  hour. 

892.  The    Basic    Bessemer    Process. — The    process    just 
described  does  not  lower  the  phosphorus  content  of  the  pig 
iron  appreciably.     Since  the  finished  steel  must  not  contain 
more  than  a  small  fraction  of  one  per  cent  of  phosphorus,  pig 
iron,  which  carries  more  than  the  upper  limit  of  phosphorus 
allowed  in  steel,  is  unsuitable  for  the  Bessemer  process  in  its 
original  form.     While  a  large  part  of  the  pig  iron  produced  in 
the  United  States  is  sufficiently  low  in  phosphorus,  that  pro- 
duced in  Germany  runs  over  2  per  cent  of  this  element.     Two 
Englishmen,  Thomas  and  Gilchrist,  modified  Bessemer 's  process 
by  lining  the  converter  with  basic  materials,  such  as  burnt 
magnesite  (largely  MgO)  and  by  adding  limestone  along  with 
the  charge.     The  strongly  basic  slag  produced  takes  up  the 
phosphorus  of  the  pig  iron,  forming  calcium  phosphate,  and 
gives  steel  low  in  phosphorus.     In  the  basic  Bessemer  process,  a 
large  part  of  the  heat  is  derived  from  the  burning  of  the  phos- 
phorus.    To  be  suitable  for  this  process  the  pig  iron  must  con- 
tain 2  per  cent  or  more  of  phosphorus.     The  basic  process 
(Thomas-Gilchrist  process)    is  not  much  used  in  America. 

893.  The  Open  Hearth  Process. — Progress  is  ever  the  watch- 
word of  the  successful  chemical  technologist.     If  a  given  proc- 
ess has  some  inherent  defect  or  limitation,  he  tries  to  invent 
a  new  one  to  overcome  the  difficulty.     Now,   the  Bessemer 


6O2 


Introduction  to  General  Chemistry 


HEARTHv 


process  is  suitable  for  pig  iron  with  less  than  o.i  per  cent  of 
phosphorus  (acid  lining)  or  more  than  2  per  cent  (basic  lining) 
but  not  for  material  of  intermediate  composition.  The  pud- 
dling (pig  boiling)  process  (891)  is  handicapped  because  it 
is  limited  to  the  production  of  batches  of  a  few  hundred  pounds 
and  because  the  temperature  is  not  high  enough  to  keep  the 
charge  melted  to  the  end  of  the  purification  and  thus  allow 
the  production  of  steel.  The  open  hearth  process  surmounts 
all  these  difficulties.  It  is  one  that  combines  the  advantages 
of  the  puddling  and  Bessemer  processes  and  produces  high 
grade  steel  from  material  having  a  wide  range  of  composition 
with  respect  to  phosphorus  and  other  components  as  well.  The 
development  of  this  process  was  made  possible  by  the  inven- 
tion by  Sir  William  Siemens, 
an  Englishman,  of  a  method  of 
producing  an  extremely  high 
furnace  temperature  by  the  use 
of  gaseous  fuel  and  without  re- 
course to  the  heat  furnished  by 
the  oxidation  of  the  silicon, 
phosphorus,  or  carbon  of  the 
charge.  Siemens'  invention 
consists  in  a  method  of  heating 
the  fuel  gas  and  air  separately 

to  a  high  temperature  before  the  gas  is  burned.  This  is 
accomplished  by  the  use  of  the  waste  heat  of  the  very  hot 
gaseous  products  of  combustion  that  come  from  the  fur- 
nace. The  process  resembles  somewhat  that  by  which  the 
air  supplied  to  the  blast  furnace  is  heated  in  stoves  (887). 
In  this  case,  however,  both  fuel  gas  and  air  are  preheated. 
Since  both  fuel  and  air  are  at  a  high  temperature  before  they 
combine,  the  heat  of  combustion  raises  the  temperature  to  a 
point  well  above  the  melting  point  of  pure  iron,  1530°.  This 
temperature  is  about  400°  higher  than  the  melting  point  of  pig 
iron.  The  essential  features  of  the  Siemens  open  hearth  fur- 
nace are  shown  diagrammatically  in  Fig.  126.  The  fuel  gas 
burns  in  the  space  above  the  hearth,  the  waste  gases  pass  out 


FIG.  126 


Metallurgy  603 

and  traverse  one  of  the  two  brick  checker  work  preheaters  from 
which  they  pass  up  the  chimney.  At  the  same  time,  air  and 
fuel  gas  are  being  heated  separately  by  passage  through  the 
other  preheater  which  previously  has  been  heated  white 
hot  by  waste  gases.  Every  half  hour  the  direction  of  flow  of 
the  waste  gases  on  the  one  hand  and  the  fuel  gas  and  air  on  the 
other  is  changed  so  that  one  preheater  is  being  heated  while 
the  other  is  giving  up  its  heat  to  the  fuel  and  air. 

The  open  hearth  process  may  be  used  with  either  acid 
(silica)  or  basic  (magnesia)  lining  just  as  in  the  Bessemer  proc- 
ess. But  it  is  the  basic  process  that  is  of  far  the  greater  im- 
portance, since  this  allows  the  treatment  of  charges  with  all 
ranges  of  phosphorus  content.  Furthermore,  by  a  procedure 
introduced  by  Messrs.  Martin  of  France,  scrap  steel  is  usually 
added  to  the  pig  iron  charge;  hence  arises  the  name,  the 
Siemens-Martin  open  hearth  process.  At  the  present  time 
about  three-fourths  of  all  the  steel  produced  in  the  United  States 
is  made  by  this  process. 

The  process  is  carried  out  by  charging  the  furnace  (basic  lin- 
ing) with  20  to  60  tons  of  pig  iron  (either  cold  or  molten),  scrap 
steel,  and  scrap  iron,  together  with  several  tons  of  lime  or  lime- 
stone. As  the  charge  melts  down,  part  of  the  silicon  and  man- 
ganese are  oxidized  and  pass  into  the  slag.  The  addition  of 
iron  ore  at  this  stage  accomplishes  further  oxidation  and  elimina- 
tion of  all  important  impurities,  excepting  sulfur,  which  is  only 
partly  removed.  The  regulation  of  the  temperature,  the  com- 
position of  the  slag  and  the  intensity  of  oxidation  are  all  sub- 
ject to  easy  control.  The  products,  metal  and  slag,  remain 
molten  until  a  chemical  analysis1  of  the  steel  shows  the  correct 
composition.  The  metal  is  then  cast  into  ingots. 

894.  Other  Processes  for  Making  Steel. — The  cementation 
process  of  making  steel  antedates  the  other  process  already 
described.  It  consists  in  heating  wrought  iron  in  contact  with 
powdered  charcoal  to  a  high  temperature  for  several  days.  The 
iron  strips  or  bars  are  packed  in  the  charcoal  powder  and  con-" 

1  Every  steel  works  has  a  chemical  laboratory  and  a  staff  of  chemists.  A  steel 
analysis  can  be  completed  in  about  20  minutes. 


604  Introduction  to  General  Chemistry 

tained  in  closed  iron  boxes  which  are  then  heated  in  a  furnace. 
Sufficient  carbon  is  taken  up  by  the  iron  to  change  it  into  a 
product  which,  with  proper  subsequent  heat  treatment,  be- 
comes extremely  hard  and  elastic.  This  steel  was  used  for 
knives,  tools,  etc.  The  product  of  the  cementation  process  was 
of  variable  composition,  owing  to  unequal  absorption  of  carbon 
even  by  the  different  iron  bars  of  a  single  lot.  A  greatly  im- 
proved and  very  uniform  product  is  produced  by  melting  cemen- 
tation steel  in  crucibles  of  clay  or  graphite.  The  product,  called 
crucible  steel,  is  the  highest  grade  steel  made.  It  is  used  for 
razors,  dies,  fine  steel  instruments,  etc. 

During  the  last  15  or  20  years  electric  arc  furnaces  have 
come  into  rather  extensive  use  for  the  manufacture  of  steel. 
These  furnaces  allow  far  easier  control  of  conditions  such  as 
temperature,  oxidizing,  or  reducing  atmosphere,  acidity  or  ba- 
sicity of  flux,  etc.  than  is  possible  in  any  other  process.  So- 
called  electric  steel  seems  to  offer  great  possibilities. 

895.  The  Working  of  Iron  and  Steel.- — In  our  description  of 
the  Bessemer  and  open  hearth  processes  we  ended  in  each  case 
with  the  casting  of  ingots,  pieces  of  metal  weighing  several  Ions. 
As  a  rule  these  are  not  allowed  to  cool  but  are  worked  up  into 
merchantable  shapes,  such  as  rails,  beams,  plates,  etc.,  in  the 
rolling  mills  that  form  an  important  part  of  every  steel  works. 
The  ingots  while  still  red  hot  are  drawn  from  the  molds  and 
carried  at  once  to  a  furnace  in  which  they  are  heated  uniformly 
until  the  metal  is  somewhat  plastic;  since  fortunately  iron  has 
the  property  of  becoming  gradually  softer  as  its  temperature  is 
raised.  The  white  hot  ingot  is  then  run  through  a  series  of 
power  driven  rolls  having  grooved  faces,  for  the  production  of 
the  desired  finished  shapes.  Each  pair  of  rolls  diminishes  the 
cross  section  of  the  piece  and  increases  its  length.  Finally  the 
finished  pieces  are  cut  into  the  desired  lengths,  while  still  hot, 
by  steel  saws. 

Many  shapes  that  cannot  be  produced  by  rolling  are  made 
by  forging  the  hot  metal  by  the  aid  of  huge  power  driven  ham- 
mers. Complex  shapes  are  made  by  the  use  of  steel  dies  into 
which  the  hot  plastic  metal  is  forced  by  powerful  hammer  blows. 


Metallurgy  605 

The  welding  of  iron  is  a  process  of  uniting  two  pieces  of  iron 
by  hammering  or  pressing  them  together  at  a  white  heat.  In 
welding,  the  surfaces  to  be  united  should  be  clean  and  free 
from  oxide.  Borax  is  usually  used  as  a  flux  to  dissolve  the 
oxide  (803).  Only  wrought  iron  and  mild  (low  carbon)  steel 
can  be  welded.  The  use  of  the  oxy acetylene  torch  in  the  cut- 
ting and  welding  of  iron  has  already  been  mentioned  (315). 
In  acetylene  welding,  a  rod  of  nearly  pure  iron  is  used  to  furnish 
molten  iron  to  weld  together  two  adjacent  pieces  of  metal. 

896.  The  Effect  of  Sulfur  and  Phosphorus  in  Iron  or  Steel.— 
Sulfur  is  always  an  objectionable  impurity  in  iron  and  steel  and 
its  amount  should  never  exceed  0.05  of  one  per  cent.     If  present 
in  larger  proportions,  sulfur  causes  cracks  to  form  in  the  metal, 
especially  when  the  latter  is  worked  at  a  red  heat.     The  be- 
havior of  sulfur  seems  to  be  due  to  the  formation  of  easily  fusible 
ferrous  sulfide    FeS  (339,  601).     The  presence  of  manganese 
tends  to  counteract  the  bad  effect  of  sulfur  by  forming  with  it  a 
sulfide,  MnS,  which  is  less  harmful  than  FeS. 

The  presence  of  phosphorus  in  steel  in  amounts  over  o.i  of 
one  per  cent  causes  marked  brittleness  in  the  cold  steel.  The 
danger  is  increased  by  reason  of  the  fact  that  as  the  steel 
solidifies  in  the  ingot  mold  (891),  the  phosphorus  compounds 
tend  to  segregate  in  the  last  portions  to  solidify  (upper  central 
portion).  If  this  part  of  the  ingot  is  not  rejected,  the  rails  (for 
example)  will  be  defective  in  places  where  the  phosphorus  con- 
tent is  high.  The  failure  of  such  high  phosphorus  rails  has  led 
in  the  past  to  numerous  railroad  wrecks.  These  facts  were 
established  in  numerous  cases  by  chemical  analysis  of  the  rails 
that  had  failed. 

The  presence  of  phosphorus  in  cast  iron  is  in  some  cases 
desirable,  and  may  be  as  high  as  one-half  of  one  per  cent,  or 
sometimes  higher.  Phosphorus  increases  the  fluidity  of  molten 
cast  iron  and  makes  easier  the  production  of  thin  castings.  In 
small  proportions,  it  is  said  to  increase  the  strength  of  the  cast 
iron. 

897.  The  Tempering  of  Steel.— The  most  important  differ- 
ence between  steel  and  iron  is  the  quality  of  the  former  by  rea- 


606  Introduction  to  General  Chemistry 

son  of  which  it  may  be  rendered  hard  or  soft  by  certain  heat 
treatments.  If  steel  containing  about  0.9  per  cent  of  carbon  is 
suddenly  cooled  from  a  bright  red  heat  (about  800°)  to  ordinary 
temperature  it  is  rendered  extremely  hard.  At  the  same  time 
its  elasticity  and  brittleness  are  increased.  If  the  same  sample 
of  hardened  steel,  after  being  again  heated  red  hot,  is  cooled 
very  slowly  it  will  be  made  so  soft  that  it  can  easily  be  filed 
or  cut  with  a  lathe  tool.  Files  and  lathe  tools  are  made  of 
hardened  steel.  It  is  usually  desirable  to  produce  steel  having 
less  than  the  maximum  hardness  and  brittleness.  This  is  ac- 
complished by  partially  drawing  the  hardness  of  completely 
hardened  steel  by  heating  it  for  a  short  time  to  200  to  300°  and 
then  suddenly  cooling  the  piece  with  water.  This  procedure 
is  called  tempering.  In  this  way,  it  is  possible  to  produce  a  prod- 
uct of  the  desired  hardness  and  toughness.  Iron  containing 
little  or  no  carbon  cannot  be  hardened  by  sudden  cooling. 

For  various  classes  of  steel,  the  ordinary  practical  limits  of 
carbon  lie  between  0.2  and  1.5  per  cent;  since  with  greater 
proportions  of  carbon  the  hardened  steel  is  too  brittle  to  be 
serviceable. 

The  causes  of  the  variations  of  properties  of  steel  and  cast  iron 
with  percentage  composition  and  with  heat  treatment  have 
been  the  subject  of  very  extensive  investigations,  for  which 
neither  money  nor  brains  have  been  lacking.  A  great  deal  has 
been  accomplished;  but  much  still  remains  to  be  discovered. 
The  matter  is  so  complex  that  it  can  only  be  touched  upon  here. 
Iron,  like  sulfur  (600),  phosphorus  (582),  carbon  (630),  and 
many  other  elements  can  exist  in  more  than  one  allotropic 
form.  The  three  forms  of  iron  are  known  as  alpha  (a)  iron,  beta 
(|8)  iron  and  gamma  (7)  iron.  The  first,  a,  is  stable  at  ordinary 
temperatures  and  changes  to  the  second,  /3,  at  750°.  Upon 
being  cooled,  /3  iron  changes  to  a  iron  below  this  temperature. 
The  case  is  like  that  of  rhombic  and  monoclinic  sulfur,  for  which 
the  transition  temperature  is  95°  (600). 

Beta  iron  is  stable  between  750°  and  860°  at  which  latter 
temperature  it  changes  into  7  iron.  The  two  transition  points 
for  a-0  and  (3-y  iron  may  be  found  experimentally  by  appli- 


Metallurgy  607 

cation  of  the  same  principle  that  enables  one  to  find  the  melting 
point  of  ice.  When  ice  melts,  heat  is  absorbed  (118).  A  ther- 
mometer surrounded  by  ice  in  a  warm  room  remains  at  o°  until 
the  ice  is  all  changed  to  water.  Now  when  a  iron  changes  to 
ft  iron  (at  750°)  heat  is  absorbed.  In  consequence,  if  a  block  of 
iron  is  placed  in  a  furnace  at  1000°,  for  example,  its  temperature 
will  rise  steadily  up  to  750°  and  then  remain  stationary  at  this 
point  until  the  change  of  a  to  /3  is  completed.  The  temperature 
of  the  iron  block  will  again  rise  steadily  until  it  reaches  860°, 
where  it  stops  once  more  during  the  change  of  0  to  7.  Upon 
cooling  the  reverse  occurs:  the  fall  of  temperature  halts  at 
two  places  corresponding  to  the  two  transition  points. 

Let  us  now  turn  to  another  topic,  the  carbon  content  of  the 
iron.  Iron,  like  many  other  metals  (calcium,  silicon,  aluminum, 
631)  forms  a  carbide.  This  has  the  formula  Fe3C  and  contains 
6.7  per  cent  of  carbon.  This  carbide  of  iron  is  known  as 
cementite.  It  decomposes  into  its  constituents  above  1000°. 

Pure  iron  melts  at  1530°.  At  this  temperature  it  dissolves 
several  per  cent  of  carbon.  1)he  solidification  point  (freezing 
point)  of  iron  is  lowered  by  dissolved  carbon  just  as  is  the 
freezing  point  of  water  containing  a  dissolved  substance  (718). 
The  maximum  lowering  is  produced  by  4.3  per  cent  of  carbon. 
Iron  of  this  composition  (cast  iron)  solidifies  at  1140°.  At  the 
same  time,  more  than  half  of  the  4.3  per  cent  of  carbon  separates 
out  as  graphite,  which  gives  to  gray  iron  its  characteristic  ap- 
pearance. The  remainder  of  the  carbon,  about  2.0  per  cent, 
remains  uniformly  disseminated  through  the  solid  iron  (gamma 
form)  and  is  said  to  form  a  solid  solution  called  martensite  in 
the  7  iron.  At  temperatures  below  1000°  this  carbon  of  the 
solid  solution  can  unite  with  iron  to  form  the  carbide  Fe3C 
(cementite) . 

Now,  it  appears  that  the  changes  of  7  into  0  iron  and  0  into  a 
iron  are  greatly  hindered  by  the  presence  of  carbon,  so  that  if 
steel  containing,  let  us  say,  i.o  per  cent  of  carbon  is  cooled 
quickly  to  ordinary  temperature,  the  change  to  a  iron  does  not 
occur  to  an  appreciable  extent.  The  very  hard  and  brittle 
product  is  therefore  largely  the  solid  solution  of  carbon  in  0 


608  Introduction  to  General  Chemistry 

(or  7)  iron,  called  martensite.  When  this  is  tempered  by  gentle 
heating  to  200  to  300°,  more  or  less  of  the  martensite  changes  to 
a  iron  (ferrite)  and  cementite,  whereby  the  product  increases 
in  toughness  and  decreases  somewhat  in  hardness. 

The  most  useful  method  of  studying  the  physical  composition 
of  steel  is  by  means  of  the  microscopic  examination  of  polished 
and  etched  samples.  Photomicrographs  made  of  such  samples 
always  show  steel  to  be  of  heterogeneous  structure.  The  study 
of  alloys  in  this  fashion  is  called  metallography.  The  metallo- 
graphic  examination  of  steel  forms  an  indispensable  part  of  the 
tests  made  in  practice  to  determine  whether  the  material  comes 
up  to  the  specifications  of  the  contract  on  which  it  is  bought. 

898.  Alloy  Steels. — When  steel  tools  are  heated  and  then 
cooled  slowly,  the  steel  loses  its  "temper"  and  they  become 
worthless.    As  a  consequence,  the  invention  of  special   alloy 
steels  which  are  indifferent  to  heat  treatment  has  been  of  tre- 
mendous importance  for  cutting  tools  used  on  lathes.     These 
are  tungsten  steels  which  have  been  subjected  to  a  special  heat 
treatment.     With  the  best  carbon-steel  tool  a  lathe  could  be 
run  at  about  30  feet  a  minute  without  spoiling  the  tool.     With 
a  high  speed  steel  tool  it  can  be  run  at  the  rate  of  300  feet  a 
minute.     The  tool  may  even  become  red  hot  without  losing  its 
cutting  power. 

Many  other  steel  alloys  are  in  use  for  special  purposes.  For 
example  a  14  per  cent  manganese  steel  is  used  for  burglar  proof 
safes,  since  it  is  too  hard  to  be  cut  or  broken;  silicon-manganese 
steels  are  used  for  springs;  nickel-chromium  steel  for  engine 
shafts,  projectile  cases,  and  armor  plate.  The  familiar  "  stain- 
less steel"  used  in  cutlery  is  a  chromium  steel  containing  small 
proportions  of  cobalt,  manganese,  etc.  It  is  not  stained  by 
contact  with  food  stuffs.  Vanadium  steel  is  much  used  for 
automobile  parts  because  of  its  great  tensile  strength  and 
elasticity. 

899.  Copper  and  Its  Ores. — Copper  ranks  next  to  iron  in 
economic   importance.     Over  half   of   the   world's   output   of 
copper  is  produced  in  the  United  States.     Four  states,  Arizona, 
Montana,  Utah  and  Michigan  produce  over  80  per  cent  of  the 


Metallurgy  609 

domestic  copper.  Much  of  the  ore  is  converted  into  metallic 
copper  in  the  states  where  it  is  mined.  A  great  variety  of 
copper  bearing  minerals  are  known;  but  the  number  of  practical 
importance  is  small. 

Metallic  or  native  copper  is  the  chief  ore  in  Michigan.  In 
the  western  states,  sulfide  ores  predominate,  the  most  impor- 
tant being  chalcopyrite,  CuFeS2.  Ores  of  copper  containing 
arsenic,  antimony,  iron,  lead,  zinc,  silver,  gold  or  other  elements 
also  occur.  Near  the  earth's  surface,  in  place  of  sulfides  there 
are  found  oxides,  carbonates  (e.g.  malachite,  34),  silicates,  etc., 
which  have  formed  by  the  natural  oxidation  of  the  sulfides  and 
reaction  with  water,  carbon  dioxide,  silica,  etc. 

The  minimum  copper  content  that  an  ore  must  have  to  be  of 
commercial  value  depends  on  several  factors,  among  which 
are  the  locality  of  the  deposit,  the  cost  of  extracting  the  metal, 
and  the  prevailing  price  of  metallic  copper.  At  present  (1919) 
the  Michigan  mines  producing  native  copper  can  be  run  at  a 
profit  on  ores  carrying  less  than  one  per  cent  of  this  metal. 
The  average  copper  content  of  ores  being  mined  at  present  in 
the  United  States  is  less  than  2  per  cent. 

900.  Ore  Dressing. — If  an  ore  contains  but  2  per  cent  of 
copper  it  will  usually  consist  to  a  large  extent  of  rock  and  earthy 
material,  the  so-called  gangue.  The  gangue  of  copper  ores  is 
usually  high  in  silica  and  silicates  of  aluminum,  magnesia, 
calcium,  etc.  Sometimes  carbonates  of  the  last  two  elements 
are  present.  Ores  from  the  western  states  usually  contain  a 
large  percentage  of  pyrite,  FeS2.  Most  copper  ores  before 
being  smelted  are  put  through  mechanical  processes  which 
separate  a  large  part  of  the  gangue  from  the  ore  as  mined 
and  thus  give  concentrates  much  richer  in  copper  than  the 
original  ores.  These  processes  are  known  in  general  as  ore 
dressing.  The  possibility  of  concentrating  ores  like  those 
here  considered  rests  on  several  radically  different  principles. 
The  most  important  of  these,  on  which  all  of  the  older  methods 
of  concentration  are  bases,  depends  the  fact  that  the  valuable 
minerals  are  all  much  denser  than  the  gangue.  Thus  the 
density  of  chalcopyrite  is  4.2  while  that  of  quartz  (silica)  is 


6io  Introduction  to  General  Chemistry 

2.6.  In  concentrating  ores  the  latter  are  crushed  so  as  to 
produce  small  fragments  which  consist  either  of  fairly  pure 
mineral  or  of  worthless  gangue.  The  fine  material  is  then 
screened  or  otherwise  treated  so  as  to  classify  the  product  with 
respect  to  size  of  grains.  If  now  grains  of  approximately 
uniform  size  are  agitated  with  water,  the  denser  grains  of 
mineral  will  tend  to  settle  to  the  bottom  of  the  container  so 
that  the  gangue  can  be  washed  away  with  the  water.  Numer- 
ous ingenious  devices  for  the  practical  application  of  this 
principle  are  in  extensive  use.  These  comprise  jigs,  spitz  - 
kasten,  vanners,  Wilfley  tables,  etc.,  descriptions  of  which  are 
to  be  found  in  most  books  on  metallurgy.  The  miner's  pan 
by  which  gold  was  freed  from  gangue  by  washing  with  water 
was  the  original  concentrating  device. 

The  flotation  process  of  concentration  consists  in  violently 
agitating  finely  powdered  sulfide  ores  with  water  and  air  and 
a  minute  quantity  of  an  oil.  The  oil  adheres  to  sulfide  minerals 
as  a  thin  film  and  the  oil  coated  particles  become  attached  to 
minute  air  bubbles  and  are  thus  carried  to  the  surface  of  the 
water  as  a  froth  while  the  worthless  gangue  settles  to  the  bottom. 
This  process  has  in  recent  years  become  of  very  great  importance 
in  the  concentration  of  the  ores  of  copper,  zinc,  lead,  etc. 

By  means  of  various  methods  of  dressing  it  is  usually  practi- 
cable to  eliminate  two- thirds  or  more  of  the  worthless  gangue. 

901.  Principles  of  the  Metallurgy  of  Copper. — The  principal 
problem  confronting  the  metallurgist  of  western  America  is 
the  preparation  of  copper  of  99.95  per  cent  purity  or  better 
from  ores  containing  less  than  2  per  cent  of  this  metal  together 
with  far  larger  proportions  of  iron,  sulfur  and  silica,  and  consid- 
erable amounts  of  many  or  all  of  the  elements,  lead,  zinc, 
aluminum,  calcium,  magnesium,  arsenic,  antimony,  and  small 
amounts  of  many  others,  notably  selenium,  tellurium,  bismuth, 
silver  and  gold.  In  the  solution  of  this  problem  he  must  not 
only  make  pure  copper;  but  he  must  get  a  good  recovery  of 
the  ore  content  of  this  metal  and  at  the  same  time  save  nearly 
the  whole  of  the  valuable  silver  and  gold  present.  It  is  easy 
to  see  that  this  problem  is  far  more  complex  than  that  involved 


Metallurgy  611 

in  the  production  of  pig  iron  (888).  Naturally  the  details  of 
the  metallurgical  processes  will  vary  in  different  localities  by 
reason  of  considerable  variation  in  the  nature  of  the  ores. 
Nevertheless  the  general  procedure  is  tolerably  uniform  through- 
out the  whole  of  western  America,  including  Canada  and 
Mexico.  The  usual  process  consists  of  three  or  four  separate 
steps.  In  the  four-step  process  the  ores  are  first  strongly 
heated  in  air  (roasted)  to  oxidize  iron,  sulfur  and  arsenic  and 
drive  off  the  last  two  as  oxides.  The  product  (calcine)  is 
then  melted  down  (smelted)  in  huge  reverberatory  furnaces 
(890)  to  form  two  products,  slag  and  matte,  both  liquid  at  a 
white  heat.  The  matte  is  largely  a  crude  sulfide  of  copper 
and  iron  and  contains  also  the  silver  and  gold  of  the  ore.  In 
the  third  step  the  molten  matte  is  oxidized  by  a  blast  of  air  in  a 
vessel  resembling  a  Bessemer  converter.  Here  the  sulfur  is 
eliminated  as  862,  the  iron  oxidized  and  combined  with  the 
silica  of  the  lining  of  the  converter  to  form  a  slag,  and  the 
copper  set  free  as  crude  metal  containing  also  silver  and  gold. 
The  fourth  and  final  step  in  the  purification  is  the  electrolytic 
refining  of  the  crude  copper.  In  the  three  step  process,  the 
first  two  steps  above  mentioned  are  combined.  In  this  case 
the  ore,  rich  in  sulfur  (pyrite),  is  smelted  directly  in  a  blast 
furnace  somewhat  resembling  that  used  for  iron  (889).  Here 
a  large  part  of  the  heat  is  supplied  by  the  burning  of  the  pyrite 
(FeSa),  the  remainder  being  furnished  by  the  burning  of  coal 
added  to  the  charge.  The  non-gaseous  products  are  again  slag 
and  matte.  The  latter  is  finished  as  above  described. 

It  must  be  evident  that  a  great  variety  of  chemical  reactions 
take  place  in  the  processes  here  outlined.  We  shall  consider  the 
more  important  ones  in  connection  with  the  detailed  descrip- 
tion that  follows. 

902.  The  Roasting  Process. — A  diagram  of  the  vertical 
section  of  a  common  type  of  roasting  furnace  is  shown  in  Fig. 
127.  This  is  a  cylindrical  structure  15  feet  or  more  in  diameter 
having  6  or  7  hearths,  one  above  the  other.  The  heating 
furnace  is  at  the  base  and  the  ore  is  fed  in  at  the  top.  Power 
driven,  water  cooled  revolving  arms  stir  the  ore  and  on  one 


6l2 


Introduction  to  General  Chemistry 


hearth  slowly  rake  it  to  a  center  opening  where  it  falls  to  the 
next  lower  hearth.  Here  the  rakes  are  so  set  as  to  carry  the  ore 
to  the  circumference  where  it  again  falls  to  the  hearth  below. 
The  material,  usually  finely  divided  concentrates,  is  heated  to 
bright  redness;  but  not  hot  enough  to  melt  it.  Much  heat 
is  developed  by  the  oxidations  that  occur  in  the  process,  so 
that  for  ores  rich  in  pyrite  little  or  no 
other  fuel  is  required  once  the  process  is 
started. 

In    the  roasting  process  the  following  re- 
actions occur: 

FeS2+O2->FeS+SO2 


If  ore  with  about  30  per  cent  of  S  is  used, 

the  roasted  product  (calcine)  contains  7  or  8 
FIG.  127 

per  cent  S. 

903.  The  Reverberatory  Smelting  Process. — Modern  copper 
reverberatories  are  of  enormous  size,  having  hearths  20  X  140 
feet.  These  furnaces  are  heated  with  powdered  coal  blown 
in  by  an  air  blast,  in  this  way  producing  the  maximum  heating 
effect  of  the  fuel  in  the  smelting  chamber.  The  great  heat 
energy  of  the  waste  combustion  gases  is  utilized  to  produce 
steam  power  for  running  the  plant. 

In  the  reverberatory  the  ore  melts  and  its  silica  content 
unites  with  the  oxides  of  iron  and  smaller  amounts  of  aluminum, 
calcium,  magnesium,  etc.  to  form  a  slag  which  is  nearly  free 
from  copper.  Further  oxidation  of  sulfides  takes  place  with 
the  formation  of  SO2.  Any  oxides  of  copper  now  present  react 
with  iron  sulfide  to  form  cuprous  sulfide,  Cu2S  and  oxide  of 
iron,  which  latter  passes  into  the  slag  as  a  silicate.  The  sulfides 
of  copper  and  iron  and  of  other  metals  in  much  small  propor- 
tions form  at  a  white  heat  a  fluid  matte,  insoluble  in  molten 
slag.  The  former,  being  of  about  twice  the  density  of  the  latter, 
sinks  to  the  bottom  of  the  hearth  and  is  tapped  out  from  time 
to  time.  The  yield  of  matte  is  about  one-fourth  of  the  weight 


Metallurgy  613 

of  the  charge.     It  contains  35  per  cent  or  more  of  copper  and 
most  of  the  silver  and  gold. 

904.  The  Converter  Process.—  A  copper  converter  resembles 
a  Bessemer  converter  (Fig.  125);  but  in  some  instances  it  is 
much  larger  and  may  have  a  capacity  of  more  than  60  tons 
of  matte.  In  the  copper  conversion  process,  the  sulfur  and 
iron  are  oxidized  and  eliminated  almost  completely,  the  former 
as  SC>2,  the  latter  as  ferrous  silicate  (slag).  The  silica  is  supplied 
from  the  clay  and  quartz  sand  converter  lining  or,  in  more 
modern  practice,  by  the  addition  of  silica  with  the  charge  and 
the  use  of  a  converter  lined  with  calcined  magnesite,  which  is 
but  little  attacked.  The  matte  is  charged  in  the  white  hot 
molten  state  immediately  after  it  is  tapped  from  the  reverbera- 
tory.  During  the  conversion,  which  lasts  several  hours,  instead 
of  15  to  20  minutes  as  in  the  Bessemer  process,  the  temperature 
is  maintained  by  the  great  heat  of  combustion  of  the  sulfur  and 
the  iron.  In  addition  to  862,  the  products  of  the  conversion  are 
slag  (mostly  ferrous  silicate)  and  crude  metallic  copper  (blister 
copper).  The  formation  of  metallic  copper  is  largely  brought 
about  by  the  following  reactions: 


2Cu2O+Cu2S—  >6Cu+SO2 
3Cu2O+FeS-+6Cu+FeO+SO2 

905.  The  Copper  Blast  Furnace.  —  The  furnaces  recently 
constructed  by  the  Anaconda  (Mont.)  Copper  Company  are 
the  largest  blast  furnaces  ever  built  for  any  purpose.  They 
are  rectangular  in  section  4^  to  6  X  87  feet  inside  and  take 
a  charge  15  feet  deep.  These  furnaces  have  a  daily  capacity  of 
2500  tons.  The  charge  consists  of  lumps  of  unroasted  (raw) 
ore  and  a  small  proportion  of  coal  or  coke.  Powerful  blasts 
of  (unheated)  air,  introduced  through  numerous  tuyeres, 
supply  oxygen  for  the  burning  of  the  coal  and  also  a  large  part 
of  the  sulfur.  The  reactions  that  occur  are  similar  to  those 
that  take  place  in  the  processes  of  roasting  (902)  and  rever- 
beratory  smelting  (903)  already  described.  The  gaseous 
products  are  S02,  CO2  and  water  vapor.  Much  arsenic  is 


6 14  Introduction  to  General  Chemistry 

volatilized  as  As203.  This  is  largely  collected  after  it  condenses 
upon  cooling  along  with  other  valuable  dust  in  specially  de- 
vised dust  collectors.  The  liquid  products  are  slag  and  matte. 
The  latter  goes  directly  to  the  converters. 

906.  Lake  Copper. — The  native  copper  ores  of  Michigan 
occur  in  a  small  region  on  the  south  shore  of  Lake  Superior. 
For  many  years  these  deposits  were  the  most  productive  in 
America.     Michigan  now  ranks  fourth  among  the  copper  pro- 
ducing states.     The  metallurgy  of  the  Michigan  ores  is  simple, 
since  they  contain  free  copper  in  a  high  state  of  purity.     The 
ores  are  crushed  and  concentrated,  and  the  concentrate  smelted 
in  reverberatory  furnaces  with  enough  limestone  to  yield  a  good 
slag.     Since  the  copper  is  present  in  the  ore  in  metallic  form, 
it  is  merely  melted  in  this  process,  without  in  the  main  being 
altered  chemically.     Copper  from  this  district  is  known  in  trade 
as  Lake  copper. 

907.  The  Refining  of  Copper. — Crude  copper  (blister  copper) 
obtained  by   the  processes   here   described   must   always   be 
refined.     Blister  copper  is  melted  in  a  reverberatory  furnace 
in  an  oxidizing  atmosphere.     Sometimes  air  is  blown  into  the 
melted  metal  through  an  iron  pipe.     A  small  amount  of  slag 
and  dross  (consisting  of  oxides  of  various  metals)  is  skimmed 
off.     To  remove  cuprous  oxide,  Cu2O,  which  is  now  present 
(in  solution)  in  the  molten  copper,  the  latter  is  stirred  with  poles 
of  green  wood.     When  the  oxide  has  been  thus  reduced  to 
metal,  the  latter  is  tapped  from  the  furnace  and  cast  into 
anode  plates  or  other  forms.     The  metal  is  now  over  99  per 
cent  pure;  but  it  is  not  pure  enough  to  be  used  as  wire  for  elec- 
trical transmission  or  for  the  construction  of  electrical  apparatus 
such  as  generators,  motors,  transformers,  etc.,  since  the  elec- 
trical conductance  of  copper  is  very  greatly  decreased  by  the 
presence     of    minute    amounts    of    impurities.     Furthermore 
all  western  copper  contains  silver  and  gold  in  amounts  sometimes 
representing  20  per  cent  or  more  of  the  value  of  the  product. 
Very   complete   removal   of   impurities   and   recovery   of   the 
precious  metals  is  accomplished  by  a  process  of  electrolytic 
refining.     The  process  is  very  simple  in  principle.     If  a  solu- 


Metallurgy  615 

tion  of  copper  sulfate  is  electrolyzed  with  copper  electrodes, 
the  anode  passes  into  solution  as  Cu+  +  ions.  These  migrate 
toward  the  cathode  while  SO  4  ions  travel  in  the  opposite 
direction.  Cupric  ions  arriving  at  the  cathode  take  up  electrons 
(487)  and  deposit  in  metallic  form.  The  sulfate  ions  remain 
in  solution.  In  practice,  the  cast  anodes  (referred  to  above) 
are  about  2^  X  3  feet  and  over  an  inch  thick^at  the  start 
while  the  cathodes  are  extremely  thin  sheets  of  very  pure  copper. 
Alternate  cathodes  and  anodes  are  hung  in  a  solution  of  copper 
sulfate  and  free  sulfuric  acid  contained  in  a  lead  lined  wooden 
vat.  The  anodes  are  all  connected  to  the  positive  wire  from  the 
electric  generator,  the  cathodes  to  the  negative  wire.  The  rate 
of  electrolysis  follows  Faraday's  law  (403).  About  i  ounce 
of  copper  is  deposited  in  twenty-four  hours  for  each  ampere 
of  current.  During  the  electrolysis  the  impurities  either 
pass  into  solution  and  remain  dissolved  (arsenic,  nickel,  iron, 
zinc)  or  fall  to  the  bottom  of  the  vat  as  insoluble  residue  (silver, 
gold,  lead,  etc.).  The  electrolyte  is  purified  from  time  to  time 
and  the  mud-like  residue  worked  up  for  silver  and  gold.  The 
cathodes  are  removed  when  they  have  reached  a  thickness  of 
half  an  inch,  me  tied  down  and  cast  into  ingots,  bars,  rods, 
etc.  The  metal  now  contains  but  a  few  hundred ths  of  i 
per  cent  impurity. 

908.  Uses  of  Copper. — The  various  electrical  industries 
consume  the  bulk  of  the  copper  produced.  Next  to  silver, 
copper  is  the  best  electrical  conductor.  Copper  is  used  ex- 
tensively in  making  various  utensils  (kettles,  stills,  condenser 
worms,  etc.).  Much  copper  is  used  in  the  making  of  brass, 
an  alloy  containing  about  70  per  cent  of  copper  and  30  per 
cent  of  zinc.  Brass  is  superior  to  copper  for  many  purposes. 
It  can  be  readily  cast  and  easily  machined  and  is  much  harder 
and  more  elastic  than  copper.  Brass  is  cheaper  than  copper 
because  zinc  is  much  the  cheaper  of  the  two  components. 
Many  other  alloys  of  copper  are  in  use.  Bronze  contains 
copper,  tin,  zinc,  and  sometimes  lead.  German  silver  contains 
copper,  zinc  and  nickel.  All  United  States  coins  contain  copper: 
one-cent  pieces  contain  95  per  cent  copper,  3  per  cent  of  tin 


616  Introduction  to  General  Chemistry 

and  2  per  cent  of  zinc.  Nickel  five-cent  pieces  (nickels) 
contains  75  per  cent  of  copper  and  25  per  cent  of  nickel. 
Gold  and  silver  coins  contain  10  per  cent  of  copper. 

909.  The  Occurrence  and  Production  of  Lead.  —  In  1913, 
the  last  year  prior  to  the  war,  the  world's  production  of  lead 
was  about  1,300,000  tons.     The  United  States  was  the  largest 
producer,  furnishing  about  one-third  of   the  whole.     During 
the  war  the  United  States  production  increased  considerably. 
In  1913,  92  per  cent  of  the  domestic  production  came  from 
four  states,  which,  with  the  percentages  produced  by  each, 
were:  Missouri,  35;  Idaho,  31;  Utah,  17;  Colorado,   10.     The 
principal  ore  is  galena,  PbS,  although  the  carbonate,  cerussite, 
PbCO3   and   the  sulfate  anglesite,  PbSO4,  are  of  some  im- 
portance.   Lead  ores  often  contain  copper,  silver  and  gold  as 
valuable  components  and  also  zinc  sulfide,  which  is  as  a  rule 
detrimental. 

910.  Principles  of  the  Metallurgy  of  Lead.  —  Lead  is  even 
more  easily  recovered  from  its  ores  than  is  copper.     The  various 
oxides  are  all  easily  reduced  when  heated  with  hydrogen  or 
carbon  (coal  or  coke).    Lead  oxide,  PbO,  is  formed  by  heating 
the  carbonate  and  also  by  the  oxidation  of  the  sulfide  by  air 
(roasting). 


The  roasting  of  galena  also  gives  rise  to  lead  sulfate. 
PbS+2O2-+PbSO4 

At  a  high  temperature  lead  sulfide  reacts  with  the  oxide  or 
the  sulfate  to  form  the  metal  and  sulfur  dioxide. 

PbS+2PbO->3Pb-fSO2 
PbS+PbSO4-»2Pb+2SO2 

Lead  sulfide  also  reacts  with  iron  at  a  red  heat  to  form  lead  and 

ferrous  sulfide. 

PbS+Fe->Pb+FeS 

All  of  the  reactions  here  mentioned  take  place  in  the  smelting 
of  galena. 


Metallurgy  617 

911.  Lead  Smelting. — In  the  United  States  nearly  all  lead 
smelting  is  carried  out  in  blast  furnaces.     These  resemble  in 
principle  those  used  for  iron  (887)  and  copper  (905)  but  are  of 
smaller  size.     The  charge  usually  consists  of  lead  ore  and  coal 
or  coke.     The  lead  ore,  which  ma'y  have  an  average  lead  content 
of  20  per  cent,  is  in  part  roasted  to  remove  a  large  part  of 
the  sulfur.     In  the  furnace  the  charge  melts  down,  giving  in 
addition  to  gases  and  vapors  three  products  which  are  liquid  at 
a  red  heat:  lead,  slag  and  matte.     The  lead  is  set  free  by  the 
reactions  above  given.     The  slag  is  formed  from  the  silica  gangue 
of  the  lead  ore  and  the  added  limestone  and   iron  ore.     It 
consists  largely  of  silicates  of  calcium  and  iron.     The  matte 
(903),  which  consists   principally  of   ferrous  sulfide,  contains 
in  form  of  sulfide  nearly  the  whole  of  the  copper  content  of  the 
ore.    -The  slag  and  matte  are  tapped  out  together  in  ladles. 
The  matte,  being  heavier,  settles  to  the  bottom  and  is  easily 
separated  from  the  slag  when  the  two  have  solidified.     The 
copper  content  of  the  matte  is  recovered  as  described  earlier 
(904).     The  lead  which  accumulates  in  a  well  at  the  bottom 
of  the  blast  furnace  is  drawn  off  through  a  siphon  and  cast  into 
pigs,  usually  called  base  bullion. 

912.  The  Refining  of  Lead. — The  chief  impurities  of  the 
base  bullion  (crude  lead)  smelted  from  western  United  States 
ores  are  antimony,  arsenic,  copper,  bismuth,  silver  and  gold. 
Since  the  first  three  of  these  impurities  render  the  lead  hard, 
the  process  by  which  they  are  eliminated  is  termed  softening. 
The  bullion  is  melted  in  a  reverberatory  furnace  in  an  oxidizing 
atmosphere.     Part  of  the  antimony  and  arsenic  are  volatilized 
and   the  remainder   together  with   the   copper  pass  into   the 
bxidized  dross  that  is  skimmed  off.     The  further  purification 
of  the  lead,  which  also  accomplishes  the  recovery  of  the  silver 
and  gold,  is  usually  carried  out  in  either  of  two  ways.     The 
first  and  older  of  these  is  the  Parkes  process  in  which  a  small 
amount  of  zinc  is  added  to  the  molten  lead,  and  the  mixture 
well  stirred.     Upon  being  allowed  to  stand,  the  zinc,  which  has 
dissolved  most  of  the  silver  and  gold,  rises  to  the  surface  and  is 
skimmed  off.     This  process  is  in  principle  precisely  like  that 


618  Introduction  to  General  Chemistry 

by  which  ether  shaken  with  a  water  solution  of  bromine  removes 
most  of  the  latter  and  then  separates  as  a  layer  floating  on  the 
denser  water.  The  lead  which  has  thus  been  freed  from  silver 
and  gold  is  again  heated  in  a  reverberatory  furnace  to  remove 
the  zinc  and  small  amounts  of  other  metallic  impurities  that 
may  still  be  present.  It  is  then  cast  into  pigs. 

The  Betts  process  for  refining  lead  consists  in  the  electrolysis 
of  a  solution  of  lead  fluosilicate,  PbSiF6  (272),  using  heavy 
anodes  of  crude  lead  and  starting  with  very  thin  cathodes  of 
pure  lead.  The  process  resembles  in  a  general  way  the  electro- 
lytic refining  of  copper.  In  this  case  the  impurities  fall  to  the 
bottom  of  the  vessel  as  an  insoluble  mud,  while  very  pure 
lead  deposits  on  the  cathode.  Silver,  gold  and  bismuth  are 
recovered  from  the  mud. 

Lead  from  the  Missouri  district  does  not  contain  enough 
silver  (or  gold)  to  make  its  recovery  profitable. 

913.  Uses  of  Lead. — With  the  exception  of  iron,  lead  is  the 
cheapest  metal.     Its  normal,  pre-war  price  was  about  4^  cents 
per  pound.     As  metal,  alloys  and  compounds,  it  serves  a  great 
variety  of  useful  purposes.    Lead  lined  tanks  are  extensively 
used  in  chemical  works  for  sulfuric  acid  solutions.     Chambers 
built  of  sheet  lead  are  used  in  making  this  acid  (616).     Shot 
and  bullets  are  made  of  lead.     Solder,  which  consists  of  lead 
and  tin,  melts  at  a  lower  temperature  than  either  of  the  separate 
metals.     Type  metal  (811)  contains  a  large  percentage  of  lead. 
Fusible  alloys   (812)   also  usually  contain  lead.     Two  of  the 
most  important  uses  of  lead  are  found  in  the  manufacture  of 
"white  lead"  (basic  lead  carbonate)  which  is  extensively  used 
in  making  paint,  and  in  the  construction  of  so-called  storage 
batteries. 

914.  The  Storage  Battery.— A  storage  cell  has  electrodes  of 
lead  immersed  in  20  per  cent  sulfuric  acid.     When  the  cell  is 
charged,  one  of  the  electrodes  is  coated  with  a  layer  of  lead 
dioxide,  PbC>2  (167,  326),  the  other  with  a  deposit  of  metallic 
lead   in   spongy   form.     Upon   closing   the   circuit   a    current 
passes  through  the  connecting  wire  in  a  direction  indicating 


Metallurgy  619 

a  flow  of  electrons  from  the  lead  covered  plate  to  that  holding 
the  lead  dioxide  (471-2). 

At  the  same  time  the  following  reaction  takes  place  in  the  cell : 

PbC2+2H2SO4-hPb->2PbSO4+2H2O. 

This  cell  may  most  simply  be  considered  an  oxidation-reduction 
cell,  in  which  the  lead  dioxide  is  the  oxidizing  agent,  and  the 
spongy  lead  the  reducing  agent  (502-4).  Each  atom  of  the 
spongy  lead  passing  into  solution  as  a  Pb+  f  ion  leaves  behind  on 
the  electrode  two  electrons.  Each  molecule  of  lead  dioxide  tends 
to  yield  Pb+  "  and  2O~~.  But  the  tetravalent  lead  ion 
being  a  powerful  oxidizing  agent  easily  takes  up  from  its  sup- 
porting electrode  two  electrons  and  is  thereby  reduced  to  Pb+  +  . 
The  excess  of  electrons  set  free  at  the  spongy  lead  (negative) 
electrode  flows  through  the  wire  to  the  lead  dioxide  (positive) 
electrode  where  there  is  a  deficiency  of  electrons,  thus  consti- 
tuting the  current  produced  by  the  cell.  The  Pb + +  ions  formed 
unite  almost  completely  with  864  ~'~  ions  of  the  acid  to  form 
nearly  insoluble  lead  sulfate  which  deposits  on  the  two  elec- 
trodes. The  O~  ~  ions  of  the  Pb02  unite  practically  completely 
with  H+  ions  of  the  acid  to  form  water. 

The  storage  cell  is  so  called  because  it  is  readily  recharged 
by  passing  a  current  (from  a  power  circuit)  through  it  in  reverse 
direction  to  the  normal.  In  the  charging  process  the  changes 
described  all  take  place  in  the  reverse  sense.  The  extent  to 
which  a  cell  is  charged  is  indicated  by  the  density  of  the  sulf uric 
acid;  since  when  the  charge  is  complete,  that  is,  when  all  lead 
sulfate  has  been  changed  to  metallic  lead,  lead  dioxide  and 
sulfuric  acid,  the  density  of  the  solution  is  at  a  maximum. 
On  th^  other  hand  when  the  active  materials  are  completely 
changed  to  insoluble  lead  sulfate  and  water,  the  cell  is  discharged 
and  the  density  of  the  liquid  is  at  a  minimum. 

The  batteries  of  electric  cars  and  the  starting  batteries  of 
gasoline  cars  are  in  general  lead  storage  cells.  Such  batteries 
deteriorate  if  they  remain  uncharged  for  a  long  time,  owing  to 
the  fact  that  the  lead  sulfate  hardens  in  a  compact  mass  which 
i§  not  easily  acted  upon  by  the  current.  It  is  therefore  advisable 


620  Introduction  to  General  Chemistry 

to  make  occasional  measurements  of  the  density  of  the  battery 
fluid  (by  means  of  a  hydrometer)  to  ascertain  when  the  battery 
needs  recharging. 

It  may  be  well  to  point  out  here  that  a  storage  cell  does  not 
store  electricity.  It  stores  chemical  energy.  The  formation 
of  two  formula  weights  of  lead  sulfate,  as  shown  by  the  fore- 
going equation,  liberates  87,000  calories  (see  Chap.  XVI)  if 
the  reaction  takes  place  directly  (as  in  a  beaker).  In  the  cell 
this  energy  is  changed  into  electrical  energy  (508),  no  appreci- 
able heat  being  produced.  When  the  cell  is  charged,  electrical 
energy  is  changed  into  chemical  energy  (509). 

915.  The  Occurrence  and  Production  of  Zinc. — Zinc  (tech- 
nically called  spelter,  148)  ranks  third  in  point  of  tonnage  and 
value  among  the  metallic  products  of  the  United  States,  being 
preceded  by  iron  and  copper  and  closely  followed  by  lead. 
(881).  The  United  States  production,  which  amounted  to  but 
16,000  tons  per  annum  in  1875,  reached  685,000  tons  in  1917. 
In  1913,  with  a  production  of  about  half  that  of  1917,  the 
United  States  made  over  30  per  cent  of  the  world's  output 
of  zinc.  In  1918  the  domestic  proportion  was  much  larger. 

Over  85  per  cent  the  domestic  zinc  ores  came  from  six 
districts;  these  with  their  respective  shares  of  the  total  output 
were  as  follows:  Joplin  district  (Missouri  and  near-by  parts 
of  Oklahoma,  Kansas  and  Arkansas),  30  per  cent;  Montana, 
20;  Oklahoma-Kansas  district,  15;  Colorado,  8;  Wisconsin, 
8;  Idaho,  5.  The  remainder  came  from  a  number  of  other 
states,  including  New  Jersey,  Utah,  Nevada,  California  and 
Tennessee. 

The  principal  ore  is  sphalerite,  the  sulfide,  ZnS,  popularly 
called  blende  or  black-j£ck.  Smithsonite,  the  carbonate, 
ZnC03;  willemite,  a  silicate,  Zn2SiO4;  and  franklinite,  a  com- 
plex ore  containing  oxides  of  zinc,  iron  and  manganese  are  also 
of  some  importance. 

Ore  to  be  commercially  valuable  must  contain  a  large 
percentage  of  zinc.  In  general,  some  process  of  concentration 
(ore  dressing,  900)  is  used  to  bring  the  zinc  content  up  to  the 
required  percentage.  In  many  cases  sphalerite  (blend)  is 


Metallurgy  621 

associated  with  galena,  PbS,  from  which  it  can  also  be  separated 
mechanically  (900). 

916.  Principles  of  the  Metallurgy  of  Zinc. — The  first  step 
in  the  treatment  of  either  carbonate  or  sulfide  ores  is  the  con- 
version of  these  compounds  into  oxide.     In  the  case  of  the 
carbonate   this  is  readily   accomplished  by  heating  the  ore, 
thus  causing  the  elimination  of  carbon  dioxide. 

ZnC03-»ZnO+C02 

The  sulfide  is  changed  to  oxide  by  heating  it  to  a  red  heat  with 
free  access  of  air  (roasting,  902) . 

2ZnS+3O2->2ZnO+2SO2 

This  process  of  oxidation  may  also  go  further  and  yield  zinc 
sulfate  as  the  result  of  the  catalytic  oxidation  of  sulfur  dioxide 
to  trioxide  (615)  and  the  union  of  the  latter  with  zinc  oxide. 
The  roasting  of  blende  is  carried  out  somewhat  like  that  of 
sulfide  ores  of  copper  (902).  Or  by  a  newer  process,  the  oxida- 
tion may  be  made  more  rapid  by  use  of  an  air  blast  so  that  a 
higher  temperature  is  reached  and  the  product  is  partly  fused 
so  that  the  small  particles  are  bound  together  in  larger  porous 
lumps,  a  process  called  sintering.  Roasting  reduces  the  sulfur 
content  of  the  ore  or  concentrate  from  20  per  cent  or  more  to 
about  i  per  cent. 

Zinc   oxide   is   readily  reduced  to  metal  by  the  action  of 
carbon  at  a  high  temperature. 

ZnO  +  C  -» Zn  +  CO  (329). 

Since  the  boiling  point  of  zinc,  918°,  is  not  far  from  the  tem- 
per^ture  at  which  this  reaction  takes  place,  the  zinc  is  volatilized 
at  the  same  time  and  thus  freed  almost  completely  from  im- 
purities. The  vapors  upon  cooling  deposit  metallic  zinc. 

917.  Details  of  the  Smelting  Process. — The  calcined  car- 
bonate or  roasted  sulfide  ore  is  mixed  with  powdered  coal  and 
placed  in  specially  constructed  fire  clay  retorts.     The  latter  are 
4  or  5  feet  in  length  and  8  or  10  inches  in  diameter,  and  have 
an  opening  at  one  end  through  which  the  gaseous  products 
escape.    The  retorts  are  set  in  a  specially  constructed  furnace, 


622 


Introduction  to  General  Chemistry 


RCTORfS 


FORHACB 


Fig.  128,  in  several  superimposed  tiers.  Each  retort  is  fitted 
with  an  attached  cylindrical  pipe  which  extends  beyond  the 
furnace  and  in  which  the  zinc  vapors  cool  sufficiently  to  con- 
dense. During  the  reaction  in  the  retort,  carbon  monoxide 
escapes  and  burns  at  the  outer  end  of  the  condenser.  A  small 
part  of  the  zinc  separates  in  the  condenser  in  powdered  form, 
known  as  zinc  dust.  This  is  a  mixture  of  metal  and  oxide, 
which  is  usually  added  to  the  next  charge. 

Since  the  charge  of  each  retort  is  only  about  1000  pounds  of 
ore,  a  zinc  smelting  works  requires  several  thousand  retorts. 

It  would  thus  appear  that  the  art 
of  zinc  smelting  is  in  a  very  back- 
ward state  of  development  in  com- 
parison with  the  present  methods 
of  smelting  iron,  copper,  and  lead. 
In  fact,  the  method  of  zinc  smelt- 
ing here  described  (the  Belgian 
process)  has  undergone  no  essen- 
tial change  in  many  decades. 

918.  Electrolytic  Zinc.— Al- 
though the  retort  distillation  pro- 
cess is  the  one  by  which  the  bulk 
of  the  zinc  is  at  present  produced,  another  process  that  looks 
very  promising  is  coming  into  rather  extensive  use.  In  this 
process  the  zinc  content  of  the  ore  is  brought  into  water  solu- 
tion as  sulfate  and  electrolyzed  in  lead  lined  wooden  vats  with 
graphite  anodes  and  aluminum  cathodes,  on  which  the  pure 
zinc  deposits.  At  the  anode  oxygen  is  set  free  and  the  cor- 
responding amount  of  hydrogen  from  the  decomposed  water 
yields  with  the  S04 —  ions  of  the  solution  sulfuric  acid.  After 
this  zinc  has  been  removed  from  the  solution,  the  latter  is 
used  to  prepare  a  fresh  zinc  sulfate  solution. 

The  conversion  of  the  zinc  content  of  the  ore  into  sulfate  is 
carried  out  in  two  ways.  By  one  process  the  ore  is  smelted  in  a 
large  reverberatory  furnace  at  a  high  temperature  and  with  an 
oxidizing  atmosphere.  The  zinc  is  reduced  to  metal;  the 
latter  escapes  as  vapor  which  burns  at  once  to  zinc  oxide  (148). 


FIG.  128 


Metallurgy  623 

The  furnace  gases,  carrying  a  dense  cloud  of  oxide,  are  cooled 
and  passed  through  numerous  large  bag  niters  in  which  the 
white  solid  oxide  is  collected.  This  oxide  is  then  dissolved 
in  dilute  sulfuric  acid  (spent  electrolyte)  to  give  zinc  sulfate 
solution.  The  other  process  consists  in  roasting  the  ore  in  a 
furnace  somewhat  like  that  shown  in  Fig.  127,  in  such  a  way 
as  to  convert  as  much  as  possible  of  the  zinc  sulfide  into  sulfate. 
The  roasted  material  is  then  leached  with  the  dilute  sulfuric 
acid  solution  from  the  electrolytic  tanks  and  the  resulting  crude 
zinc  sulfate  solution  purified  and  later  electrolyzed. 

919.  Uses  of  Zinc. — Zinc  is  brittle  at  ordinary  temperatures, 
but  it  becomes  malleable  and  ductile  at  150°  and  may  be  rolled 
into  thin  sheets  or  drawn  into  wire.     Metallic  zinc  does  not 
oxidize  readily  on  exposure  to  air,  on  which  account  it  is  used 

-for  containers  of  various  sorts.  Far  more  important  is  the  use 
of  zinc  for  making  galvanized  iron.  This  is  iron  (sheet,  wire, 
pipe,  etc.)  cleaned  with  sulfuring  acid  and  coated  with  zinc  by 
dipping  it  in  the  molten  metal.  This  treatment  protects  the 
iron  from  rusting;  since  even  if  the  iron  is  imperfectly  covered, 
the  zinc  has  the  greater  tendency  to  pass  into  the  ionic  state 
(492)  that  is,  to  oxidize  (504).  Brass  (908)  and  other  alloys  of 
zinc  require  for  their  manufacture  immense  quantities  of  zinc. 
As  negative  electrodes  in  electric  batteries  (496)  (e.g.  dry  cells) 
zinc  finds  an  important  use.  In  the  laboratory  and  in  chemical 
works  zinc  is  used  as  a  reducing  agent  and  to  produce  hydrogen. 
Zinc  oxide  and  zinc  sulfide  are  white  solids  insoluble  in  water 
which  are  used  extensively  in  making  white  paints.  Lithopone, 
which  is  also  used  for  the  same  purpose,  is  a  mixture  of  zinc 
sulfide  and  barium  sulfate  made  by  mixing  solutions  of  barium 
sulfide  and  zinc  sulfate. 

BaS-fZnSO4-»BaSO4+ZnS 

920.  Aluminum. — Although   aluminum    (174)    is    the   most 
abundant  truly  metallic  element,  constituting  over  seven  per 
cent  of  the  earth's  crust  (F.  W.  Clarke),  it  was  also  the  latest 
metal   to  come  into  extensive  use.     The  reason  for  its  tardy 
practical  development  is  found  in  the  fact  that  none  of  the  proc- 


624  Introduction  to  General  Chemistry 

esses  that  serve  to  reduce  to  the  metallic  state  ores  of  iron, 
copper,  lead,  zinc,  tin,  mercury,  nickel,  silver  or  gold  are  effec- 
tive in  the  case  of  aluminum.  This  metal  was  probably  first 
obtained  by  Wohler  (696)  in  impure  form  in  1827  and  in  a  purer 
form  in  1845,  by  heating  aluminum  chloride  with  metallic 
potassium. 


Deville  improved  the  process  by  substituting  for  the  too  easily 
volatile  A1C13  its  double  salt  with  NaCl  and  made  it  much 
cheaper  by  using  sodium  in  place  of  the  very  expensive  potas- 
sium. It  was  through  the  efforts  of  this  great  French  chemist 
that  aluminum  began  to  be  manufactured  in  1856.  In  the 
year  1859  the  metal  sold  for  90  cents  an  ounce.1  This  high 
price,  which  was  largely  due  to  the  cost  of  the  sodium  required 
for  Deville's  process,  made  this  very  desirable  metal  too  expen- 
sive for  very  extended  use.  It  was  therefore  a  great  advance 
when  in  1886  Castner  invented  a  new  process  for  the  manufac- 
ture of  sodium  by  which  the  latter  could  be  produced  for  25 
cents  per  pound.  Castner's  process  consisted  in  reducing  caus- 
tic soda  at  a  red  heat  with  iron  and  carbon.  By  1889  the 
price  of  aluminum  made  by  use  of  Castner's  cheap  sodium  had 
dropped  to  $4.00  a  pound  ($8.80  a  kilo).  But  only  two  years 
later  aluminum,  made  by  an  entirely  different  process  which 
did  not  require  metallic  sodium,  was  put  on  the  market  in  quan- 
tity at  $i  .co  a  pound.  How  this  came  about  will  be  shown  in 
the  following  section. 

921.  The  Electrical  Production  of  Aluminum.  —  At  the  begin- 
ning of  the  1  9th  century  none  of  the  metals  of  the  alkalies  or 
the  alkaline  earths  had  been  isolated,  although  the  bases  and 
salts  of  these  elements  were  well  known.  The  preparation  of 
these  new  metals  by  Davy  at  the  Royal  Institution  (London)  in 
1807  was  the  most  important  advance  in  our  science  since  the 
discovery  of  oxygen  and  the  explanation  of  its  role  in  chemical 
change,  more  than  thirty  years  earlier.  Davy  made  sodium  by 
passing  an  electric  current  from  a  large  battery  of  galvanic  cells 

1  This  corresponds  to  $14.40  a  pound  or  $31.70  a  kilo.  In  the  year  1913  the 
price  in  the  United  States  was  21  cents  a  pound,  or  46  cents  a  kilo! 


Metallurgy  625 

through  molten  (water  free)  sodium  hydroxide.  Although  dry 
solid  sodium  hydroxide  at  room  temperature  is  a  very  poor 
conductor  of  electricity,  the  fused  base  conducts  readily.  Me- 
tallic sodium  collects  at  the  cathode  where  hydrogen  is  also  set 
free,  while  oxygen  comes  off  at  the  anode.  In  a  similar  way 
potassium  is  obtained  from  its  hydroxide.  In  the  following 
year  Davy  tried  to  make  aluminum  by  electrolysis,  but  was  not 
successful.  In  the  course  of  the  succeeding  three-quarters  of 
a  century  other  chemists  at  times  considered  the  possibility  of 
making  aluminum  by  electrolysis;  but  this  method  was  not 
practical  because  the  electric  generator  (dynamo)  was  not  yet 
developed  and  all  electric  power  was  obtained  at  great  cost  from 
batteries,  which  derived  their  energy  from  metallic  zinc. 

The  era  of  electro-chemistry  began  about  1880  with  the  de- 
velopment of  the  power-driven  dynamo  to  the  stage  where 
electric  power  was  readily  produced  at  moderate  cost.  The 
first  even  partially  successful  electrical  method  for  aluminum 
was  that  invented  in  1884  by  the  Cowles  brothers  of  Cleve- 
land. In  this  process  a  powerful  electric  current  was  passed 
through  a  mixture  of  aluminum  oxide  and  charcoal  in  the 
presence  of  metallic  copper.  At  the  extremely  high  temperature 
produced  in  the  Cowles  electric  furnace  the  aluminum  oxide 
was  reduced  by  the  charcoal  and  was  dissolved  by  the  molten 
copper  and  so  protected  from  further  change.  The  product  was 
aluminum  bronze,  an  alloy  with  copper  containing  15  to  20 
per  cent  of  aluminum.  This  was  not  an  electrolytic  process 
since  it  worked  equally  well  with  an  alternating  current.  The 
first  real  electrolytic  process  for  aluminum  was  developed  by 
Hall  of  Oberlin,  Ohio,  in  1886. 

At  nearly  the  same,  time  a  process  practically  identical  with 
Hall's  was  patented  by  Herault  in  France.  This  process,  which 
since  its  invention  has  displaced  all  others  for  the  manufacture 
of  aluminum,  is  commonly  known  as  the  Hall-Herault  process, 
to  the  description  of  which  we  may  now  turn. 

922.  The  Hall-Herault  Process. — In  order  to  make  clear 
the  principles  of  the  Hall-Herault  process,  a  few  lines  must  be 
devoted  to  the  matter  of  the  electrical  behavior  of  fused  salts. 


626  Introduction  to  General  Chemistry 

Although  dry  solid  salts  are  very  poor  electrical  conductors 
(384)  they  usually  conduct  even  better  than  most  water  solu- 
tions when  they  are  in  the  fused  molten  state.  It  seems 
reasonable  to  suppose  that  fused  salts  are  highly  ionized  (411) 
and  that  their  electrical  conductivity  is  to  be  explained  by  the 
migration  of  charged  ions  just  as  in  the  case  of  solutions.  It 
is  now  well  known  that  many  metals  (e.g.,  sodium,  potassium, 
magnesium,  calcium,  barium,  etc.)  can  be  obtained  in  metallic 
form  by  the  electrolysis  of  certain  of  their  fused  salts,  especially 
their  chlorides. 

Hall  found  that  the  electrolysis  of  the  fused  double  fluoride 
of  sodium  and  aluminum,  3NaF-AlF3,  which  is  found  in  Green- 
land as  the  mineral  cryolite,  gave  melted  metallic  aluminum. 
He  also  found  that  if  he  added  to  the  molten  cryolite  (M.P. 
1000°)  aluminum  oxide,  A^Oa  (174),  that  the  latter  dissolved 
abundantly,  giving  a  more  fusible  mixture  When  this  was 
electrolyzed  with  a  carbon  rod  anode  it  gave  aluminum  as 
before,  and,  at  the  anode,  set  free  oxygen  which  united  with  the 
carbon  to  form  carbon  dioxide.  In  this  case  since  the  oxide  of 
aluminum,  instead  of  the  cryolite,  is  decomposed,  it  is  only 
necessary  to  add  the  former  from  time  to  time  and  to  draw  off 
the  melted  metal  to  make  the  process  a  continuous  one. 

Hall's  apparatus  consists  of  a  large  iron  vessel  lined  with  a 
hard,  compact  form  of  carbon  which  serves  as  the  cathode  and 
within  which  the  molten  electrolyte  is  contained.  The  anode  is 
a  stout  rod  of  carbon  which  dips  into  the  electrolyte.  The 
strength  of  current  depends  on  the  size  of  the  cell,  specifically 
on  the  area  of  anode  surface  exposed  to  the  electrolyte.  The 
electrolysis  takes  place  at  900°  to  1000°;  the  electrolyte  is 
maintained  at  this  high  temperature  by  the  current  alone, 
without  being  heated  externally. 

923.  The  Occurrence  and  Preparation  of  Materials  for  the 
Hall  Process. — The  cryolite  needed  is  obtained  either  from  the 
natural  deposits  of  this  mineral  in  Greenland  or  it  may  be  made 
artificially  from  aluminum  hydroxide,  soda  and  hydrofluoric 
acid  (269) .  Although  aluminum  minerals  are  both  numerous 
and  abundant  everywhere,  it  has  been  found  that  only  one 


Metallurgy  627 

mineral,  bauxite,  can  be  used  as  a  technical  source  of  the  metal. 
Bauxite  (first  found  at  Beaux  in  France)  is  a  hydrated  oxide  of 
aluminum,  A^Os^H^O.  It  is  rarely  found  pure  and  contains 
usually  oxides  of  iron,  silica,  clay,  sand,  etc.  About  80  per 
cent  of  the  domestic  supply  of  bauxite  comes  from  Arkansas, 
the  remainder  coming  from  Georgia,  Alabama  and  Tennessee. 
To  prepare  the  alumina  needed  for  the  Hall  process,  bauxite 
is  digested  with  concentrated  caustic  soda  to  dissolve  the 
aluminum  as  sodium  aluminate  (177).  The  latter  gives  by- 
hydrolysis  (436)  aluminum  hydroxide,  which  in  turn  gives  the 
oxide  when  heated. 

924.  The  Production  of  Aluminum. — Since  the  invention  of 
the  Hall-Herault  process,  the  manufacture  of  aluminum  has 
increased  prodigiously.     In  1918  the  world's  production  was 
nearly  500  million  pounds,  of  which  the  United  States  made 
about  45  per  cent.     Only  four  metals,  iron,  copper,  zinc  and 
lead  are  made  in  larger  amounts  than  aluminum. 

925.  The  Uses  of  Aluminum. — Two  hundred  million  pounds 
of  aluminum  were  used  by  the  allies  in  aeroplane  manufacture 
in  1918.     Over  one- third  of  the  material  of  the  Liberty  motor 
is  aluminum.     The  automobile  industry  uses  a  large  quantity 
in  the  form  of  castings  and  in  sheets  for  chassis  and  paneling. 
Aluminum,  used  for  castings,  is  usually  alloyed  with  copper 
or  zinc  to  improve  its  workability,  etc.,  since  the  pure  metal 
shrinks  badly  in  the  mold  and  sticks  to  the  cutting  tool  when 
it   is   machined.     " Duralumin" — aluminum   and  copper  with 
traces  of  manganese  and  magnesium — has  been  much  used 
for  rolled  and  forged  parts.     Another  big  supply  of  aluminum 
is  needed  for  cooking  utensils  of  various  kinds.     It  is  interesting 
to  note  that  though  this  metal  is  more  costly  than  copper  if 
measured  by  weight,  it  is  less  costly  if  measured  by  volume. 
Hence  where  articles  of  a  certain  size  are  required  aluminum 
may  well  be  substituted  for  copper.     Aluminum  is  coming  into 
extensive  use  for  electrical  transmission  wire  and  also  for  other 
electrical  purposes.     A  wire  of  aluminum  of  the  same  weight 
and  length  as  a  copper  wire  is  twice  as  good  a  conductor  of  the 
current  (half  the  resistance). 


628  Introduction  to  General  Chemistry 

We  have  already  spoken  of  the  thermite  reaction  of  Gold- 
schmidt  (330)  in  which  iron  oxide  is  reduced  by  aluminum  at 
high  temperature  and  free  iron  and  aluminum  oxide  form.  By 
this  method  small  quantities  of  pure  iron  are  made  for  heavy 
iron  and  steel  repair  work.  The  molten  metal  formed  is  used 
to  weld  broken  parts  together.  Ihe  thermite  reaction  is  also 
employed  to  prepare  chromium,  manganese  and  vanadium 
respectively  from  their  oxides  and  also  for  the  manufacture  of 
many  ferro-alloys. 

Much  aluminum  is  also  used  in  the  steel  industry  to  prevent 
blow  holes  in  castings.  Aluminum  powder,  mixed  with  oil,  is 
used  as  a  silvery  paint. 

926.  Gold  and  Silver. — The  most  important  processes  for  the 
recovery  of  gold  from  ores  that  have  little  or  no  value  on 
account  of  other  metals  present  have  already  been  briefly 
described  (819).  The  amalgamation  process  for  gold  is  also 
suitable  for  silver  if  the  latter  is  present  in  the  free  state.  Ores 
in  which  the  silver  is  present  as  sulfide,  arsenide,  or  a  more  com- 
plex form  are  given  preliminary  treatments  of  various  sorts  be- 
fore amalgamation.  In  the  cyanide  process  for  gold,  if  silver 
is  present  it  accompanies  the  gold  to  a  greater  or  less  extent  by 
reason  of  its  ability  to  undergo  similar  reactions.  This  process, 
however,  is  primarily  one  for  the  recovery  of  gold. 

Silver  (but  not  gold)  is  recovered  from  ores  containing  no 
other  valuable  metals  by  a  process  based  on  the  conversion  of 
the  silver  present  into  chloride  by  roasting  the  ore  with  common 
salt  and  then  extracting  (leaching)  the  material  with  a  solution 
of  sodium  thiosulfate  (626).  The  silver  dissolves  as  soluble 
salt  of  a  complex  ion  and  is  recovered  by  precipitation  with 
sodium  sulfide  which  throws  down  extremely  insoluble,  black 
silver  sulfide,  Ag2S.  From  the  latter,  metallic  silver  is  easily 
made. 

Many  ores  of  gold  or  silver  or  both  also  contain  lead,  copper 
or  zinc  in  valuable  amounts.  Such  ores  are  smelted  in  the 
ways  already  described  in  this  chapter.  In  smelting  processes 
the  gold  and  silver  are  almost  completely  concentrated  in  the 
base  metals,  It  is  frequently  advantageous  to  smelt  gold  or 


Metallurgy  629 

silver  ores,  free  from  copper  or  lead,  along  with  ores  of  these 
base  metals,  especially  if  the  former  ores  contain  iron  which  is 
needed  to  make  a  suitable  slag.  About  75  per  cent  of  the  silver 
produced  in  the  United  States  is  obtained  from  the  refining  of 
copper,  lead  and  zinc. 

Since  neither  gold  nor  silver  is  oxidized  at  high  tempera- 
tures, the  refinery  residues  are  strongly  heated  under  oxidizing 
conditions  until  other  metals  have  been  converted  into  oxides 
and  thus  eliminated.  The  metallic  product,  which  contains  as 
a  rule  a  large  excess  of  silver,  is  called  dore  silver.  The  alloy 
of  the  two  metals  is  usually  treated  with  hot  sulfuric  acid  which 
converts  the  silver  into  sulfate  and  leaves  the  gold  unchanged. 
The  silver  sulfate  is  dissolved  in  water  and  reduced  to  metal  by 
the  action  of  copper  or  by  electrolysis.  The  gold  residue  is 
melted  and  cast  into  bars. 

The  uses  of  silver  and  gold  are  so  well  known  that  no  com- 
ment on  this  topic  is  necessary. 


INDEX 


INDEX 

[References  are  to  sections,  not  pages] 


Absolute,  temperature,  6 
Absorption,  731 
Acetamide,  659 

Acetates,  precipitation  by  sodium  ace- 
tate, 452 
Aceteldehyde,  652 

Acetic  acid:  as  a  dissolving  agent,  457; 
effect  of  acetates  on  ionization  of,  43 1 ; 
effect  of  hydrochloric  acid  on,  432; 
ionization  of,  409;  failure  to  precipi- 
tate acetates,  452;  graphic  formula, 
654;  general  properties,  157,  653; 
and  sodium  hydroxide,  435;  titra- 
tion  of,  440 

Acetone,  656 

Acetylene:  composition,  50;  and  copper 
oxide,  83;  heat  of  combustion,  357; 
preparation,  49;  series,  663 

Acheson  process,  630 

Acid  anhydride,  313 

Acidimetry,  137 

Acids:  and  carbonates,  163;  dibasic, 
102;  fatty,  655;  ionization  of,  410; 
monobasic,  102;  parts  of,  377;  prop- 
erties of,  90;  strength  of,  428;  tribasic, 
159;  strong:  little  soluble  salts  of, 
and  acids,  458;  and  salt  of  weak  acid, 
428,  430;  salts  of,  and  weak  acids, 
457;  titration  of,  440;  and  weak 
acid,  432;  and  weak  base,  437;  weak, 
177;  little  soluble  salts  of,  solution 
by  strong  acids,  456;  as  precipitating 
agent,  452;  and  strong  acid  on,  432; 
and  a  strong  base,  434;  suppression 
of  ionization  of,  431 ;  titration  of,  440; 
and  a  weak  base,  438 

Adhesion,  726 

Adsorption,  728,  731,  732,  763;  accom- 
panying precipitation,  739;  and 
catalysis  by  finely  divided  metals, 
73T>  by  colloid,  739;  of  gases  by 
charcoal,  728;  from  solution,  732; 
from  solution,  explanation  of,  763 

Affinity,  259 

Agate,  804,  806 

Air:  adsorption  by  glass,  729;  compo- 
sition of,  10,  765,  766;  ionization  of, 
854;  liquid,  776,  777;  weight  of,  3 

Aitken,  769 

Albumin,  685;  and  the  phosphoric 
acids,  598 


Alcohol:  absolute,  641;  denatured,  641; 
from  fermentation,  640,  641;  wood, 
645 

Alcohols:  aromatic,  670;  triatomic,  679 

Aldehydes,  652;  aromatic,  670 

Alfalfa,  assimilation  of  nitrogen,  515 

Aliphatic  compounds,  666 

Alkali  metals,  823 

Alkaline  earth  metals,  824 

Allotropic  forms,  582,  600,  630 

Alloys,  fusible,  812,  913 

Alpha  rays,  480,  858 

Alum,  175 

Aluminium.    See  Aluminum 

Aluminum,  174;  acid  reaction  of  salts 
of,  176;  metallurgy  of,  920  ff.;  place 
in  Periodic  Table,  833;  production 
of,  881,  924;  uses  of,  925;  chloride, 
174;  hydroxide:  gel,  761;  precipita- 
tion by  ammonium  hydroxide,  452; 
preparation  and  properties,  174,  177; 
nitrate,  175;  nitride,  514;  oxide,  174; 
potassium  sulfate,  175;  sodium  sul- 
fate,  175;  sulfate,  175 

Amethyst,  806 

Amides,  659 

Amine  acids,  685 

Amines,  658;  aromatic,  674 

Ammeter,  400 

Ammonia:  adsorption  by  charcoal,  728; 
composition  of,  52,  53;  and  copper 
oxide,  equation,  84;  critical  tempera- 
ture of,  775;  liquid,  517;  manufacture 
from  cyanamid,  526;  manufacture  of 
synthetic,  525;  oxidation  to  nitric 
acid,  570;  properties  of,  51,  517,  527; 
sources  of,  516;  theory  of  the  syn- 
thesis of,  520-24;  uses  of,  518;  and 
water,  91;  and  water,  equilibrium, 
284. 

Ammonium:  aluminum  sulfate,  175; 
arseno-molybdate,  810;  bicarbonate, 
dissociation  of,  530;  chloride,  92;  dis- 
sociation of  vapor,  529;  and  sodium 
hydroxide,  ionic  theory  of,  426,  462; 
cyanate,  synthesis  of  urea,  696; 
fluoride,  269;  hydrogen  sulfate,  101; 
hydroxide:  a  base,  91;  degree  of 
ionization  of,  409;  effect  of  am- 
monium salts  on  ionization,  431; 


633 


634 


Introduction  to  General  Chemistry 


[References  are  to  sections,  not  pages] 


effect  of  sodium  hydroxide  on  ioniza- 
tion, 432;  and  nitric  acid,  105; 
titration  of,  440;  molybdate,  597; 
nickel  sulfate,  817;  nitrate,  105; 
decomposition  by  heat,  565;  use  as 
an  explosive,  556,  573;  nitrite,  513; 
perchlorate,  355;  phosphomolybdate, 
597,813;  picrate,673;  sulfantimonate, 
8 1 1 ;  sulfate,  101 ;  sulfides,  607 ;  disso- 
ciation of,  530;  yellow,  607;  sulfo- 
cyanate  and  ferric  chloride  (ionic 
equilibrium),  280,  415 

Ampere,  400 

Amphoteric  substances,  177 

Analysis,  31;  by  means  of  spectroscope, 
786 

Anaxagoras,  184 

Anhydride  of  an  acid,  313 

Aniline,  674 

Animals,  dependence  on  plants,  690 

Anions,  389,  391 

Anode,  295,  389 

Anthracene,  698 

Antichlor,  611 

Antimony:  and  its  compounds,  811; 
place  in  Periodic  Table,  835,  tri- 
chloride, from  antimony  and  chlorine, 
246 

Apatite,  580 

Aqua  regia,  562 

Aquadag,  750 

Argol,  665 

Argon,  513;  in  air,  765;  discovery  and 
properties,  791;  family,  825;  molec- 
ular and  atomic  weight,  792;  place 
in  Periodic  Table,  848 

Argyrol,  750 

Aromatic  series,  666 

Arrhenius,  Svante,  405,  720 

Arsenic:  antidote  for,  810;  and  com- 
pounds, 810;  place  in  Periodic 
Table,  835 

Arsenious  sulfide,  810;  colloidal  sus- 
pension, 734,  736-38 

Arsine,  810 

Atmosphere,  765-99;  a  disperse  system, 

774 

Atomic-molecular  hypothesis,  208 
Atomic  numbers,  848,  878,  880.    See 

flyleaf  at  back  of  book 
Atomic  structure,  880 
Atomic  volume,  823 
Atomic  .weights,  216;    determined  by 

law  of  Dulong  and  Petit,  229;  oxygen 

basis,  223;  of  radioactive  substances, 

869;  relation  to  symbol  weight,  216; 

table,  800.    See  also  inside  of  back 

cover  of  book 


Atoms:  Greek  conception  of,  184;  mod- 
ern conception  of,  185;  number  in  a 
molecule,  214,  215;  relative  weights 
of,  208-13;  structure  of,  470,  880 

Attraction,  molecular,  202,  726 

Atwater,  687 

Avogadro's  hypothesis  (or  law),  193; 
application  to  suspensions  in  liquids, 
707;  applications,  210;  exactness  of, 
226 

Avogadro-van't  Hoff  hypothesis,  716 

Azote,  512;   (nitrogen)  15 

Babbitt  metal,  811 

Baking  powder:  phosphate,  593;  tar- 
trate,  665 

Balloons:   helium,  797;   hydrogen,  304 

Barium,  824;  flame,  784;  in  Periodic 
Table,  832;  chloride,  164;  peroxide, 
310,  319;  salts,  reaction  with  sul- 
fates,  380;  sulfate,  164 

Base:  diacid,  146;  ionization  of,  410; 
monacid,  146;  parts  of,  378;  strength 
of,  429;  strong:  and  salt  of  weak 
base,  429-30;  titration  of,  440;  and 
weak  base,  432;  weak,  176;  and 
strong  acids,  437;  and  strong  base, 
43 1 ;  suppression  of  ionization  of,  43 1 ; 
titration  of,  440;  and  weak  acid,  438 

Basic  nitrates,  564 

Batteries,  electric,  496 

Baume,  scale  of  specific  gravity,  618 

Bauxite,  923 

Be,  618 

Bead  test,  metaphosphate,  506;  borax, 
803 

Becquerel,  852;  rays,  853 

Beets,  sugar,  682 

Bell  metal,  809 

Benedict,  687 

Benzaldehyde,  670,  697 

Benzene,  667;  ring,  668 

Benzene  sulfonic  acid,  672,  697 

Benzoic  acid,  671 

Benzyl  alcohol,  670 

Beri-beri,  689 

Beryllium:  and  compounds,  826;  in 
Periodic  Table,  832 

Bessemer  process,  891-92 

Beta  rays,  860 

Betts  process,  912 

Birkland-Eyde  process,  manufacture  of 
nitric  acid,  568 

Bismuth:  compounds,  812;  place  in 
Periodic  Table,  835;  nitrate,  564; 
sub  nitrate,  564 

Bivalent,  146 

Blast  furnace,  886 


Index 


635 


[References  are  to  sections,  not  pages] 


Bleaching,  by  chlorine,  249 

Bleaching  powder,  351 

Blende,  915 

Blotting  paper,  functioning  of,  726 

Blowpipe,  oxyhydrogen,  300 

Bluestone,  497 

Body:  living,  need  of  food,  686;  use  of 
word  in  chemistry,  2  2 

Boiling-point:  molar  elevation,  718;  of 
solutions,  128;  of  water,  112 

Boisbaudran,  845 

Bonds:  double,  661;  triple,  663;  va- 
lence, 323 

Bone  ash,  158 

Bone  black,  use  as  adsorption  agent,  732 

Boracic  acid,  801 

Borax,  801,  803 

Boric  acid,  801,  802 

Boron:  flame,  784;  in  Periodic  Table, 
833;  properties,  80 1 

Boyle's  Law,  4;  explanation  of,  188 

Bragg,  W.  H.,  879 

Bragg,  W.  L.,  879 

Brandt,  577 

Brass,  148,  908,  919 

Breakfast,  example,  688 

Bredig,  744 

Erin's  process,  for  oxygen,  310 

Britannia  metal,  811 

British  Thermal  Unit,  358 

Bromic  acid,  822 

Bromides:  action  of  chlorine  on,  259; 
insoluble,  257 

Bromine:  liquid,  255;  occurrence  of, 
254;  in  Periodic  Table,  83  7 ;  prepara- 
tion and  properties,  255,  258;  uses 
of,  260;  water,  255 

Bronze,  809 

Brownian  movements:  of  particles  in 
liquids,  705;  of  smoke  particles,  703 

Brownlee  apparatus,  44 

B.T.U.,  358 

Bullion,  base,  912 

Bunsen  burner,  781 

Burning  of  substances,,  10-20 

Burns:  from  nitric  acid,  104;  by  steam, 
117;  from  sulfuric  acid,  93 

Butyric  acid,  680 

Butyrin,  680 

Cadmium,  832 

Caesium,  823;  most  active  metal,  842; 
in  Periodic  Table,  831 

Calcium,  150,  832;  flame,  784;  spec- 
trum of,  786;  bicarbonate,  156; 
carbide,  49,  631;  carbonate,  150; 
dissolved  by  hydrochloric  acid,  456, 
461;  precipitation  of,  448;  chloride, 


151;  action  with  carbonic  acid,  449; 
cyanamid,  514;  fluoride,  267;  hy- 
droxide, 151,  1 66;  hypochlorite,  351; 
nitride,  514;  oxide,  150;  phosphate, 
158;  precipitation  of,  452;  silicate, 
270;  sulfate,  153;  tungstate,  814 

Calomel,  182 

Calorie,  in 

Calorimeter,  bomb,  357 

Candle,  burning  of,  20,  780 

Cane  sugar,  682 

Caprilic  acid,  680 

Caproic  acid,  680 

Carat,  819 

Carbides,  631 

Carbohydrates,  682 

Carbolic  acid,  672 

Carbon:  compounds,  629-701 ;  "fixed," 
in  fuel,  359;  formation  in  flame,  782; 
heat  of  combustion,  357;  and  hydro- 
gen, 631 ;  place  in  Periodic  Table,  826, 
834;  properties  of,  630,  631;  as 
reducing  agent,  328;  valence,  648; 
weight  in  one  liter  of  gaseous  com- 
pounds, 60;  bisulfide  (or  disulfide), 
546  and  631;  dioxide,  19,  30,  633; 
adsorption  by  charcoal,  728;  from 
calcium  carbonate,  150;  composition 
of,  39;  constancy  of  concentration  in 
air,  766;  critical  temperature  of,  775; 
and  lime  water,  151;  solid,  as  a  refri- 
gerant, 633,  775;  and  water,  equilib- 
rium, 285;  disulfide,  631;  monoxide, 
632,  665;  critical  temperature  of,  775; 
heat  of  combustion,  357;  liquefaction 
of,  777;  reducing  agent,  329;  oxy- 
chloride,  695;  tetrachloride,  644 

Carbonates  and  acids,  163 

Carbonic  acid,  152;  action  with  calcium 
chloride,  449;  equilibrium,  carbon 
dioxide,  and  water,  285;  ionization 
of,  409;  and  sodium  hydroxide,  161 

Carborundum,  631 

Carnotite,  874 

Caro's  acid,  628 

Castner's  sodium  process,  920 

Catalysis  by  metals  and  adsorption,  731 

Catalytic  agent,  or  catalyzer,  239;  for 
ammonia  reaction,  522;  for  decom- 
position of  hydrogen  peroxide,  320; 
manganese  dioxide  in  preparation  of 
oxygen,  306;  manufacture  of  sul- 
furic acid,  615;  platinum^303 

Cathion,  389,  390 

Cathode,  295,389;  rays,  475,  478 

Caustic  soda,  absorbs  carbon  dioxide, 
19.  See  Sodium  hydroxide 

Cavendish,  292,  544 


636 


Introduction  to  General  Chemistry 


[References  are  to  sections,  not  pages] 


Cells,  galvanic,  496 

Celluloid,  694 

Cellulose,  684;  products,  694 

Cementation  process,  894 

Cementite,  897 

Cerium,  820,  834 

Cerium  oxide,  use  of,  in  gas  mantles,  301 

Chalcopyrite,  899 

Chamber  acid,  616 

Chamber  process,  616 

Chaptal,  572 

Charcoal:  adsorption  by,  728,  732; 
burning  of,  18 

Charge  on  a  suspension,  test  for,  749 

Charles's  Law,  5 

Chemical  activity  of  elements,  order  of, 
in  Periodic  Table,  842 

Chemical  reactions,  rate  of,  275-78 

Chile  saltpeter,  104;  iodine  from,  261 

Chlorates,  preparation  of,  353 

Chloric  acid,  354;  and  lead  sulfide,  354 

Chlorides,  insoluble,  252 

Chlorination  process  for  gold,  819 

Chlorine:  and  aluminum,  174;  and 
bromides,  259;  critical  temperature, 
775;  discovery  of,  234;  from  elec- 
trolysis of  hydrochloric  acid,  43; 
electrolytic  preparation,  237,  238; 
and  ferrous  chloride,  173;  and  hydro- 
gen, 44,  243;  liquefaction  of,  242; 
and  mercury,  179;  and  methane,  644; 
minimum  weight,  63;  occurrence, 
233;  oxidation  products  of,  352;  an 
oxidizing  agent,  332;  in  the  Periodic 
Table,  827;  and  phosphorus,  247; 
physical  property  or,  241;  poisonous 
gas,  236;  preparation  from  hydro- 
chloric acid,  235,  239;  and  sodium, 
theory  of  union  of,  485;  and  tur- 
pentine, 248;  union  with  metals,  246; 
and  uses  of,  249;  and  water,  at 
ordinary  temperatures,  245;  at  high 
temperatures,  240;  dioxide,  354; 
monoxide,  352 

Chloroform,  644 

Chlorophyl,  691 

Chloropicrin,  695 

Chromates,  345;  as  oxidizing  agents, 
346 

Chromium:  and  its  compounds,  344; 
place  in  the  Periodic  Table,  836; 
preparation  of,  925;  hydroxide,  pre- 
cipitation of,  452 

Cinnamic  aldehyde,  697 

Cinnamon,  oil  of,  697 

Citric  acid,  665 

Clark,  F.  W.,  884,  920 

Claude,  776 


Clay,  177,  804 

Cleveite,  helium  from,  794,  867 

Clouds,  768-71;  functions  of  gaseous 
ions  in  their  formation,  771 

Clover,  assimilation  of  nitrogen,  515 

Cloves,  oil  of,  697 

Coal,  distillation  of,  634 

Coal  tar,  667 

Cobalt:  and  its  compounds,  816,  817; 
place  in  Periodic  Table,  838 

Coefficient,  in  chemical  equations,  76 

Cohesion,  726 

Coins,  composition  of,  908 

Coke,  630;  as  a  reducing  agent,  328 

Collodion,  694 

Colloids,  735 ;  adsorption  of  agent 
which  precipitates,  739;  chemistry, 
importance  of,  764;  protecting  agent 
for  suspensions,  745;  test  for  charge 
on,  749 

Columbium,  835 

Combining  volumes,  Gay  Lussac's  Law, 

22O,  221 

Combustion:  heat  of,  356-58;  in  nitric 
oxide,  546;  spontaneous,  364.  See 
also  Burning  of  substances 

Common  ion  law,  432 

Complex  ions,  538 

Composition,  law  of  definite,  46,  99 

Compound,  31 

Concentration  and  speed  of  reaction, 
280 

Conductivity:  change  during  neutraliza- 
tion, 423;  effect  of  dilution,  406; 
molecular,  407 

Conservation  of  energy:  for  bodily 
processes,  687;  law  of,  371 

Conservation  of  matter,  21 

Constant  heat  summation,  law  of,  363 

Constant  of  equilibrium  (example),  283 

Constant  proportion,  law  of,  46,  99 

Contact  process  for  sulfuric  acid,  617 

Cooledge,  814 

Copper,  165;  bead  test  for,  783;  blister, 
907;  burning  0^32;  electrolytic,  907 ; 
flame,  784;  from  hydrogen  and  copper 
oxide,  33;  metallurgy,  900  ff.;  and 
nitric  acid,  550,  561;  production,  881, 
899;  properties  and  place  in  Periodic 
Table,  831;  refining  of,  907;  uses  of, 
908;  (see  also  Cupric  and  Cuprous); 
ammonium  ion,  538;  chloride,  165; 
from  chlorine  and  copper,  246; 
hydroxide,  165;  nitrate,  165;  oxide, 
165;  and  acetylene,  equation,  83; 
and  ammonia,  equation,  84;  composi- 
tion of,  32,  38;  and  hydrogen,  33,  82; 
sulfate,  165 


Index 


637 


[References  are  to  sections,  not  pages] 


Cordite,  693 

Corn  sirup,  639 

Corrosive  sublimate,  178 

Cotton,  684;  bleaching  of,  351;  soluble, 

694 

Coulomb,  400 
Cowles  brothers,  921 
Cream  of  tartar,  665 
Critical    pressure,    775;    temperature, 

775 

Crookes,  477,  859;  tube,  475 

Cryolite,  267,  923 

Crystals:  melting  of,  205;  theory  of 
structure,  204;  X-ray  and  structure 
of,  879 

Cupric  compounds  (see  also  Copper 
compounds),  333;  bromide,  color 
of,  in  solution,  396;  chloride,  elec- 
trolysis of  solution,  386;  as  catalyst, 
239;  oxide,  325;  potassium  chloride, 
175;  sulfate,  electrolysis  of  solution, 
386;  sulfide,  precipitated  by  hydro- 
gen sulfide,  452 

Cuprous  compounds,  333;   oxide,  325 

Curie,  Mme,  855 

Current,  electrical:  by  chemical  action, 
493-95;  direction  of,  472;  electron 
theory,  469;  how  carried  through  a 
wire,  471;  nature  of,  468;  strength 
of,  400 

Cyanates,  665 

Cyanide  process  for  gold,  819 

Cyanides,  665 

Dalton:  Atomic  Hypothesis,  208;  Law 

of  Partial  Pressure,  192 
Daniell  cell,  496 
Davy,  234,  921 
Deacon's  process,  239 
Decomposition,  25;  of  sal  soda,  26 
Degree  of  ionization,  408-10 
Deliquescence,  130 
Democritus,  184 
Density,  no 
Depilatory,  607 
Detonator,  573 
Developers,  photographic,  manufacture 

related  to  dye  industry,  699 
Deville,  920 
Dew,  767 

Dewar,  777,  779;  vessels,  777 
Dew  point,  767 
Dextrose,  639 
Dialysis,  743 
Dialyzer,  743 

Diamonds,  630;  and  gamma  rays,  86 1 
Diastase,  682 
Diatoms  (infusorial  earth),  732 


Dichlor  diethyl  sulfide  (mustard  gas), 

695 

Bichromates,  345 ;  oxidizing  agents,  346 

Dietetics,  688 

Diffusion:  of  gases,  191;  in  liquids, 
704 

Dinner,  example,  688 

Disintegration  hypothesis,  481,  865 

Disperse  systems,  725-64,  774 

Displacement:  electronic  interpreta- 
tion of,  489;  metallic,  electronic 
interpretation  of,  491;  of  metals  by 
one  another,  490;  of  non-metals  by 
one  another,  488 

Dissociation:  electrolytic  (see  Ioniza- 
tion); hydrolytic,  436;  of  volatilized 
ammonium  salts,  529-31 

Distillation,  23 

Dithionic  acid,  628 

Double-decomposition,  337,  383;  and 
electrical  conductivity,  384;  and  the 
ionic  hypothesis,  413 

Drying  agents,  130 

Dulong  and  Petit,  law  of,  229-30 

Duralumin,  925 

Durion,  539,  804 

Dust:  in  the  air,  768;  counting  of 
particles,  769;  explosions,  365;  func- 
tion of,  in  cloud  formation,  768 

Dynamite,  692;  preparation,  of  726 

Effervescence,  163 

Efflorescence,  131 

Egg  white,  685 

Ekaboron,  845 

Ekaluminium,  845 

Ekamanganese,  845 

Ekasilicon,  845 

Electric  cells,  oxidation-reduction,  502 

Electric  current,  468,  469;  oxidation 
and  reduction  by  means  of,  507 

Electricity :  f  rictional,  474 ;  nature  of,  466 

Electrochemical  equivalents,  law  of,  403 

Electrodes,  potential  difference  of,  499 

Electrolysis,  27;  of  cupric  chloride 
solution,  386;  of  cupric  sulfate  solu- 
tion, 386;  Faraday's  laws  of,  399; 
of  hydrochloric  acid,  385;  of  silver 
nitrate,  387;  of  sodium  chloride  solu- 
tion, 385;  terms  used  in,  389;  theory 
of,  398,  487;  of  water,  27 

Electrolytes,  389;  molecular  weights 
of,  720;  precipitation  of,  by  common 
ion,  453;  soluble,  equilibrium  be- 
tween, 441;  solution  of  little  soluble, 

455 

Electromotive  force,  499 
Electromotive  series  of  metals,  492,  499 


638 


Introduction  to  General  Chemistry 


[References  are  to  sections,  not  pages] 


Electrons,  465,  466;  mass  of,  479; 
proof  of  existence,  467;  vibration, 
cause  of  light,  788 

Element,  31,  35;  meaning  of  term, 
873;  total  number,  800 

Elements,  activity  of,  842 

Emery,  174;  powder  graded  by  time 
of  settling  in  water,  733 

Emulsifying  agents,  755 

Emulsoids,  735,  75i~57;  general  prop- 
erties, 756;  importance,  757 

Energy:  chemical,  372;  chemical,  con- 
version into  electrical  energy,  508; 
conversion  of,  371;  electrical,  500; 
electrical,  conversion  into  chemical 
energy,  509;  forms  of,  372;  kinetic, 
of  molecules,  198 

Enthodermic  changes,  366 

Enzymes,  682 

Epsom  salts,  144 

Equations,  chemical,  76;  balancing  of, 
86;  balancing  of  oxidation  and  reduc- 
tion, 561;  meaning  of,  77,  85;  prob- 
lems of,  79,  87;  review  table,  86 

Equilibrium  (general):  chemical,  274; 
constant  (example),  283;  criterion  of, 
282;  effect  of  changes  of  concentra- 
tion on,  280;  effect  of  pressure  on 
hydrogen  and  nitrogen,  523;  effect 
of  pressure  on  system  in,  287;  effect 
of  removing  one  product  of  the  re- 
action, 289;  between  electrolytes  (see 
Equilibrium  betweeen  electrolytes); 
and  heat  production,  367;  hydrogen 
and  nitrogen,  effect  of  temperature 
on,  288,  521;  kinetic  hypothesis 
applied  to,  279;  liquid  and  vapor, 
201 ;  between  molecules  and  ions,  405 ; 
physical,  273 

Equilibrium  between  electrolytes:  and 
gas  evolution,  459;  graphic  repre- 
sentation of,  417,  418  ff.,  444;  be- 
tween soluble  electrolytes,  441;  in 
solution,  413-39;  in  solution  and  solid 
substance,  443  ff. 

Equivalent,  electrochemical,  403 

Equivalent  weight,  403 

Esters,  657;  of  glycerine,  679 

Ethane,  643 

Ether,  642 

Ethyl  compounds:   acetate,  657;    alco- 
hol, 641;    structural  formula,  649; 
ammonium  iodide,  658;   iodide,  660 
Ethylene,  660;   series,  662;    structural 
formula  of,  66 1 

Ethylene  chloride,  660 

Evaporation,  theory  of,  199 
Exothermic  changes,  366 


Explosion,  302 
Explosives,  571,  692,  693 

Fajans'  Law,  872 

Falk,  722 

Families:  of  elements,  822;  A  and  B,  831 

Faraday,  242,  389,  403,  775;  law  of 
electrolysis,  399 

Fat  soluble  A,  689 

Fats,  677;  composition  of,  680 

Fehling's  solution,  683 

Fermentation,  640 

Ferric:  chloride,  and  ammonium  sul- 
focyanate,  280,  415;  and  hydrogen 
sulfide,  503;  and  potassium  iodide, 
503;  preparation,  173;  compounds, 
331;  hydroxide,  173;  colloidal,  741- 
43;  a  gel,  761;  and  hydrochloric 
acid,  455;  precipitation  of,  451; 
oxide,  173;  sulfate,  173;  and  sodium 
carbonate,  384 

Ferro-molybdenum,  813 

Ferro-silicon,  804 

Ferrortungsten,  814 

Ferrous  compounds,  331;  ammonium 
sulfate,  175;  chloride,  173;  hy- 
droxide, 173;  oxide,  173;  sulfate, 
173;  sulfide:  dissolved  by  hydro- 
chloric acid,  456;  precipitation  of, 
452 ;  preparation  and  use,  339, 601 

Fertilizer:  phosphate,  160;  use  of 
ammonium  salts  in,  518 

Films,  photographic,  694 

Filters,  adsorption  by,  763 

Filtration,  23 

Fire  extinguisher,  633,  644 

Firefly,  584 

Flame,  780;  colored,  784;  reactions  in, 
783;  spectra,  786 

Flint,  806 

Flotation  process,  900 

Flour,  bleaching  of,  563 

Fluorides,  269 

Fluorine:  most  active  non-metal,  842; 
in  Periodic  Table,  837;  preparation 
and  properties,  267 

Fluor-spar,  267 

Fluo silicates,  272 

Food,  676 

Formaldehyde,  652 

Formalin,  652 

Formic  acid,  665 

Formula:  calculation  of,  80,  81;  chemi- 
cal, 62;  of  elementary  gases,  75,  218; 
graphic,  323;  involatile  substances, 
72 ;  making  of,  67 ;  structural,  of  ethyl 
alcohol,  649;  structural,  importance 
of,  in  organic  chemistry,  651;  struc- 


Index 


639 


[References  are  to  sections,  not  pages] 


tural,  of  methyl  ether,  649;  use  of,  68; 

volatile,  liquids  and  solids,  71;  weight, 

relation   to   molecular   weight,  217; 

weight,  and  symbol  weight,  74 
Fox  fire,  584 
Franklinite,  915 
Frasch  process,  602 
Fraunhofer  lines,  790 
Frazier,  712 

Freezing-point,  molar  depression,  718 
Frost,  767 
Fruit  sugar,  682 
Fuel:    composition  of,  calorific  power, 

359;  for  steam  production,  360 
Fullers  earth,  732 

Gallium,  discovery  of,  845 

Galvanic  cells,  498 

Galvanized  iron,  148 

Gamboge,  Brownian  movements  of,  705 

Gamma  rays,  86 1 

Gas:  calorific  power,  359;  collected  over 
water,  calculation  of  pressure,  113; 
diffusion,  191;  evolution  of,  factors 
governing,  463 ;  evolution  of,  and  ionic 
equilibrium,  459;  illuminating,  634; 
natural,  643;  pressure,  cause  of,  188; 
pressure,  law  of  partial,  192;  statistics, 
194-97 

Gas  laws:  accuracy  of,  225;  applica- 
tion to  dilute  solutions,  715;  prob- 
lems on,  7 

Gases:  the  inert,  or  noble,  825;  ioniza- 
tion  of,  770;  liquefaction  of,  ^  775; 
mixing  of,  190;  standard  conditions 
for,  7 

Gasoline,  643 

Gay  Lussac's  Law,  5 

Gelatine:  gel,  758;  a  protecting  agent 
for  colloids,  745 

Gels,  758-61;  plant  and  animal  tissue, 
relation  to,  759 

German  silver,  816,  908 

Germanium:  discovery  of,  845;  place 
in  Periodic  Table,  834 

Glass:  adhesion  of  water  to,  726, 
adsorption  of  air  on  surface,  729; 
adsorption  of  water  vapor  on,  730; 
composition,  270,  808;  etching  of, 
271;  optical,  803;  quartz,  808;  vari- 
ous kinds,  808 

Glucose,  639;  fermentation  of,  640 

Glycerine,  679 

Gold:  colloidal,  746;  compounds  and 
alloys,  819;  extraction  from  ores,  819, 
926;  place  in  Periodic  Table,  831;  pro- 
duction, 88 1;  solution  in  aqua  regia, 
562;  world-production  of,  819 


Goldschmidt  process,  330,  925 

Graham,  Thomas,  735 

Grain  alcohol,  641 

Granite,  804 

Grape  sugar,  639 

Graphite,  630 

Gravity  batteries,  497 

Grindstone,  806 

Groups:  of  atoms  in  organic  chemistry, 

651;  of  elements,  830-38 
Guncotton,  693 
Gunpowder,  326,  571,  572 
Gypsum,  155 

Haemoglobin  and  oxygen,  314 

Hall-Herault  process,  921-23 

Halogens,  231,  822 

Hampson,  776 

Harkins,  754,  849 

Hartshorn,  516 

Heat:  atomic,  230;  of  combustion, 
356-58;  of  formation,  361;  ofioniza- 
tion,  439;  latent,  of  evaporation,  115; 
latent,  fusion  of  ice,  118;  law  of  con- 
stant heat  summation,  363;  mechani- 
cal equivalent  of,  369;  molecular, 
792-93;  of  neutralization,  362,  439; 
production  and  equilibrium,  367; 
production  in  physical  and  chemical 
changes,  summary  of,  366;  of  reac- 
tion, 361;  of  solution,  127;  of  solu- 
tion, 288;  of  solution  and  solubility, 
134,  288;  theory  of,  189 

Heliotrope,  697 

Helium:  and  the  alpha  rays,  795,  858; 
balloons,  797;  discovery  of,  794;  and 
family,  825;  properties  of,  796 

Hematite,  328,  884 

Henry,  Law  of,  126 

Hess,  Law  of,  363 

Hexathionic  acid,  628 

Hexoses,  682 

Hillebrand,  794 

Humidity  of  air,  766 

Hydrates,  97 

Hydrazine,  531 

Hydriodic  acid,  265,  339 

Hydrobromic  acid:  oxidation  of,  258; 
preparation  of,  256;  as  reducing 
agent,  341;  and  silver  nitrate,  257 

Hydrocarbons,  643;  aromatic,  668; 
isomerism  of,  664 

Hydrochloric  acid:  and  aluminum,  174; 
and  aluminum  hydroxide,  175;  and 
calcium  carbonate,  461;  and  caustic 
soda,  41;  and  copper  oxide,  165; 
electrolysis  of,  43,  385;  and  iron,  173; 


640 


Introduction  to  General  Chemistry 


[References  are  to  sections,  not  pages] 


and  lead,  167;  and  lead  dioxide,  157; 
and  magnesium,  149;  and  magnesium 
hydroxide,  143,  45 5  J  and  magnesium 
oxide,    145;    preparation,    103,  250; 
properties,  251, 252 ;  as  reducing  agent, 
341 ;  and  sodium  acetate  (ionic  theory 
of  action),  424;  and  sodium  hydro- 
gen sulfate,  253,  289;  and  zinc,  149 
Hydrocyanic  acid,  665 
Hydrofluoric  acid,  269 
Hydrofluosilicic  acid,  269,  272 
Hydrogen:  adsorption  by  charcoal,  728; 
critical  temperature,  775;   discovery 
of,  292;   from  electrolysis  of  water, 
27;  flame,  temperature  of,  299;  heat 
of  combustion,  357;   liquid,  779;   in 
nature,  292;  percentage  in  water,  36; 
place  in    electromotive  series,  492, 
499;    place  in  Periodic  Table,  847; 
preparation,  27,  28, 33,  293-95;  prop- 
erties of,  27,  296,  297;  use  of,  295, 
304;   weight  in  one  liter  of  gaseous 
compounds,  56;  weight  in  22.4  liters 
of  gaseous  compounds,  56,  63,  211; 
reactions  of:    burning   in   chlorine, 
244;  and  chlorine,  44;   and  copper 
oxide,  33,  82;  and  iodine,  chemical 
equilibrium,    281;  and  iodine,   heat 
of   reaction,   367;    and  iron  oxide, 
290;    from  magnesium  and  steam, 
28;  and  nitrogen,   298;  and  nitro- 
gen,   equilibrium,    521,    523;    and 
oxygen,  299-303;  chloride,  44;  com- 
position, 48;  critical  temperature  of, 
775;  physical  properties  of,  251  (see 
Hydrochloric  acid);    preparation  of, 
44;  fluoride,  268;  action  on  quartz, 
silicon,  glass,  270;   iodide:   prepara- 
tion, properties,  264;  sulfide:  aqueous, 
as  reducing  agent,  609;  precipitation 
of  sulfides,  452;  preparation  of,  prop- 
erties, 339,  605-9;   peroxide,  318-20; 
detection  of   (reaction  with  chromic 
acid),  321;  as  oxidizing  agent,  347; 
as  reducing  agent,  348;  use  of,  320 
Hydrolysis  of  salts,  436 
Hydrosulfurous  acid,  606 
Hydroxides,  preparation  of  insoluble, 

166 

Hydroxylamine,  531 
Hypo,  625,  627 
Hypochlorites,  350 
Hypochlorous  acid,  349;  from  chlorine 

monoxide,  352 
Hyposulfurous  acid,  628 

Ice:    density  of,  119;'  latent  heat  of, 

118;  manufacture  of  artificial,  5*9 
"Icy-Hot"  bottles,  777 


Ignition  temperature,  302 

Illuminating  gas,  634 

Indestructibility  of  matter,  law  of,  2 1 

Indicators,  440 

Indigo,  synthesis  of,  698 

Infusorial  earth,  732 

Ink:  sympathetic,  817;  India,  750 

Insulator,  473 

Intumescence,  803 

lodic  acid,  822;  and  sulfurous  acid,  277 

Iodides,  265;  uses  of,  266 

Iodine,  261-65;  heat  of  reaction  with 
hydrogen,  367;  and  hydrogen,  chemi- 
cal equilibrium,  281-83;  place  in 
Periodic  Table,  848;  and  starch,  263, 
637;  uses  of,  266 

Ionic  equilibrium,  405,  413  ff.;  and 
gas  evolution,  factors  governing,  463 

Ionic  hypothesis,  41 1 ;  criticism  of,  41 2 ; 
value  of,  464 

lonization:  cause  of,  486;  degree  of, 
conductivity  method,  408-10;  degree 
of,  freezing-point  method,  721,  722; 
degree  of,  graphic  representation, 
417;  of  gases,  770;  heat  of,  439 

Ions:  in  solution:  charges  on,  404;  and 
chemical  reactions,  392 ;  complex,  538; 
color  of,  395;  migration  of,  397; 
nature  of,  483;  positive  and  negative, 
393;  union  of,  394;  gaseous,  771 

Indium  and  its  compounds,  818,  838 

Iron,  173;  allotropic  forms,  897;  burn- 
ing of,  12,  17,  29,  81;  cast,  883,  889; 
economic  importance,  882 ;  galvanized, 
148,  919;  and  hydrochloric  acid,  173; 
magnetic  oxide  of,  173;  metallurgy, 
883-98;  ores,  884;  pig,  886-88;  place 
in  Periodic  Table,  838;  preparation 
of,  from  hematite,  328;  production, 
881;  and  steam,  29,  290;  and  sulf uric 
acid  (concentrated),  621;  and  sul- 
furic  acid  (dilute),  173;  welding,  895, 
925;  wrought,  883,  890 
Iron  compounds  (see  Ferrous  and 

Ferric) :  iron  oxide,  1 73 
Isatine,  698 
Isomerism,  647,  650 
Isomorphism,  Sio 
Isoprene,  700 
Isotopes,  871,  873 

Joule,  mechanical  equivalent  of  heat, 
370;  unit  of  electrical  energy,  500 

Kelly,  William,  891 
Kerosene,  643 
Ketones,  656 


Index 


641 


[References  are  to  sections,  not  pages] 


Kieselguhr,  732 

Kinetic  energy,  368;  of  molecules,  198 

Kinetic  theory,  187;  of  the  liquid  state, 

198,  704 

Kipp,  apparatus,  294 
Kraft,  577 
Krypton,  798;  and  its  family,  825 

Lactic  acid,  665 

Lactose,  682 

Lake  copper,  906 

Langmuir,  754 

Lanthanum,   820;    in   Periodic  Table, 

833,  846 
Latent  heat:  of  evaporation,  115;   of 

fusion,  118 
Laughing  gas,  556 
Laurie  acid,  680 
Lavoisier,  13 
Lead,    167;    metallurgy,   910  ff.;    place 

in  Periodic  Table,  834;   production, 

88 1,   909;    smelting,    910;    uses   of, 

913;     acetate,    167;     chloride,    167; 

dioxide,   167;    oxidizing  agent,  326; 

fluosilicate,    272,   912;     nitrate,    167; 

decomposition  by  heat,  565;    oxide, 

167;  salts,  reaction  with  sulfates,  381; 

sulfate,  167;  sulfide  and  chloric  acid, 

354;  and  hydrogen  peroxide,  347 
Lead  pencils,  630 
Leather,  artificial,  694 
Legumes,  assimilation  of  nitrogen,  515 
Levulose,  682 

Light,  ether  wave  hypothesis,  788 
Lightning,  cause  of,  773 
Lime,  slaking  of,  150 
Limelight,  301 
Limestone,  150 
Lime  water,     150;      test     for     carbon 

dioxide,  18,  151 
Liquid:    equilibrium  with  vapor,  201; 

state,  198;  supercooling  of,  206 
Litharge,  167 
Lithium,  823;  flame,  784;  spectrum  of, 

786 

Lithopone,  919 
Litmus:  reaction  to  acids,  89;  reaction 

to  bases,  88;   sensitiveness,  440 
Lubricating  oils,  643 
Luminosity,  cause  of,  in  flames,  782 
Luncheon,  example,  688 
Lye,  162 
Lynde,  776 

Madder,  698 

Magnesia,  milk  of,  142 

Magnesium:  burning  of,  n,  28,  30,  80; 

in  Periodic  Table,  832;  bromide,  255; 

chloride,  143;  fluosilicate,  272;    hy- 


droxide, 142,  143,  166;  nitrate,  145; 
oxide,  142;  sulfate,  144 

Magnetite,  884 

Malachite,  analysis  of,  34,  899 

Maltose,  641,  682 

Manganese  and  compounds,  342,  837; 
preparation  of,  925;  dioxide:  catalyst, 
306,  320;  oxidizing  agent,  326; 
preparation  of  chlorine,  234;  of 
bromine,  258 

Maple  sugar,  682 

Marble,  150 

Marine  acid,  305 

Marsh  gas,  643 

Marsh's  test,  810 

Martensite,  897 

Matches,  586 

Matte,  903,  911 

Matter:  change  of  form  with  tempera- 
ture, 9;  conservation,  law  of,  21; 
electrical  nature  of,  482,  880;  forms 
of,  2 

Meat,  685 

Medicinals,  manufacture  related  to  dye 
industry,  699 

Meker  burner,  781 

Membrane,  semipermeable,  710 

Mendelejeff,  843 

Mercuric  compounds,  333;  chloride, 
178;  nitrate,  178;  basic,  564;  oxide, 
73,  178,  181;  red  ash  of  mercury,  30; 
sulfate,  179 

Mercurous  compounds,  333;  chloride, 
180,  182;  nitrate,  180;  oxide,  181; 
sulfate,  1 80 

Mercury,  178;  ash,  decomposition  by 
heat,  14;  and  chlorine,  179;  oxides, 
181;  and  place  in  Periodic  Table,  832; 
properties,  178-82,  832;  from  red 
oxide,  14;  bichloride,  178 

Mesothorium,  870 

Metaboric  acid,  802 

Metallography,  897 

Metallurgy,  SSiff.^ 

Metals,  35;  cathions,  390;  displace- 
ment by  one  another,  490;  oxidation 
and  reduction  of,  504;  place  in 
Periodic  Table,  841;  production,  88 1 

Metaphosphoric  acid,  589,  590;  and 
albumen,  597 

Metastannic  acid,  809 

Metathesis,  383 

Methane:  analysis,  55;  properties  and 
composition,  54 

Methyl  compounds:  acetate,  657; 
alcohol,  645;  ammonium  iodide,  658; 
chloride,  644;  ether,  646;  structural 
formula  of,  649;  iodide,  660;  naph- 
thalene, 669 


642 


Introduction  to  General  Chemistry 


[References  are  to  sections,  not  pages] 


Methyl  orange,  137,  440 

Methyl  violet,  adsorption  from  solution, 

732 

Methylene  chloride,  644 

Meyer,  Lothar,  844 

Mica,  807 

Micoderma  aceti,  653 

Microcosmic  salt,  596 

Milk  sugar,  682 

Millikan,  467 

Minimum  and  multiple  weights,  ex- 
planation of  law,  212 

Minimum  weights,  63,  64,  212 

Mitscherlich,  810 

Mixed  metal,  820 

Moissan,  267,  630 

Moisture,  in  air,  766 

Molecular  heats,  792-93 

Molecular  hypothesis,  186,  208 

Molecular  weight,  217;  and  depression 
of  the  freezing-point,  718;  of  elec- 
trolytes in  solution,  720;  and  eleva- 
tion of  the  boiling-point,  718;  and 
lowering  of  vapor  pressure  of  solu- 
tion, 718;  from  osmotic  pressure, 
data,  717;  relation  to  formula  weight, 
217 

Molecules,  184, 185;  attraction  of ,  202; 
colors  of,  in  solution,  396;  kinetic 
energy  of ,  198;  motion  of,  187;  per- 
ception of,  703;  reality  of,  702; 
solubility  of,  445;  velocity  of,  197; 
velocity  of,  and  temperature,  189 

Molybdenite,  813 

Molybdenum:  and  compounds,  813; 
place  in  Periodic  Table,  836 

Molybdic  acid,  813 

Monatomic  gases,  ratio  of  molecular 
heats  in,  793 

Monazite,  820 

Mond  process  for  nickel  ores,  816 

Monel  metal,  816 

Mordant,  809 

Morse,  712 

Moseley,  848,  878 

Moth  balls,  669 

Mother  of  vinegar,  653 
Multiple  weights,  law  of,  212 
Muriatic  acid,  250 
Mustard  gas,  695 

Naphthalene,  669 

Natural  gas,  643 

Nature  of  matter,  880 

Neodymium,  820 

Neon,  798;  and  its  family,  825 

Neutrality,  of  solutions,  433 


Neutralization :  acid  and  base,  89,  3  79; 
change  of  conductivity  during,  423; 
heat  of,  362,  439;  ionic  theory 
of,  421,  423;  of  nitric  acid  by  am- 
monium hydroxide,  105;  simplified 
equation  of,  422 

Newlands,  844 

Newton,  first  law  of  motion,  187 

Nickel:  and  compounds,  816,  817; 
place  in  Periodic  Table,  838 

Nickel  ammonium  ion,  538 

Nickel  coin,  816 

Nickel  steel,  816 

Nilson,  845 

Niobium,  place  in  Periodic  Table,  835 

Niton,  825 

Nitrates:  properties  of,  564,  565; 
source  of,  540;  test  for,  549 

Nitric  acid:  from  air,  567;  and 
aluminum  hydroxide,  175;  from 
ammonia,  570;  and  ammonium 
hydroxide,  105;  anhydride,  555; 
chemical  solvent  for  salts,  560;  and 
copper  oxide,  165;  and  magnesium 
hydroxide,  145;  and  magnesium 
oxide,  145;  and  metals,  558;  from 
nitric  oxide,  569;  and  non-metals, 
559;  oxidation  by,  557;  oxidizing 
agent,  542;  properties  of,  ^104,  541; 
from  sulfuric  acid  and  niter,  104; 
test  for,  549;  uses  of,  571 

Nitric  oxide:  combustion  in,  546; 
conversion  into  nitric  acid,  569; 
critical  temperature,  775;  equilibrium 
with  nitrogen  and  oxygen,  566;  and 
•ferrous  sulfate,  548;  from  ferrous 
sulfate  and  nitric  acid,  547;  prepara- 
tion of,  543;  properties  of,  545 

Nitrides,  514 

Nitrites,  553 

Nitro  compounds,  aromatic  673 

Nitrobenzene,  673 

Nitrocellulose,  693 

Nitrogen:  assimilation  by  plants,  515; 
critical  temperature,  775;  cycle  in 
nature,  574;  discovery  of,  512; 
fixation  of  atmospheric,  575;  and 
hydrogen,  298;  and  hydrogen,  equi- 
librium, 520-25;  inert  part  of  air. 
15;  liquid,  from  liquid  air,  777; 
minimum  weight,  63 ;  occurrence  of, 
511;  and  oxygen,  544;  and  oxygen, 
equilibrium,  566;  place  in  Periodic 
Table,  835;  preparation  of,  513; 
properties  of,  514;  pentoxide,  555; 
tetroxide,  545;  physical  properties 
of,  5Si,  552;  preparation  of,  550; 
two  forms  of,  552;  trioxide,  554 


Index 


643 


[References  are  to  sections,  not  pages] 


Nitroglycerine,     692;      adsorption     in 

infusorial  earth,  726 
Nitrolime,  514 
Nitrophenol,  673 
Nitrosyl  chloride,  562,  563 
Nitrosyl  sulfuric  acid,  616 
Nitro  toluene,  ortho,  meta,  and  para, 

673 

Nitrous  acid,  553;  anhydride,  554 

Nitrous  oxide,  556;  critical  tempera- 
ture, 775 

Nonconductor  of  electricity,  473 

Non-metals,  35;  displacement  by  one 
another,  488;  oxidation  and  reduc- 
tion of,  505;  place  in  Periodic  Table, 
841 

Normal  solutions,  136-38;  595 

Noyes,  A.  A.,  722,  737 

Nutrition,  688 

Oil:  animal  and  vegetable,  657;  of 
bitter  almonds,  670,  697;  essential, 
697;  hardening  of,  68 1;  hydrogena- 
tion  of,  68 1 ;  mineral,  643;  of  vitriol, 

93 

Oildag,  750 
Oleic  acid,  680 
Olein,  680 
Oleum,  613 
Onyx,  806 
Opal,  804,  806 
Open  hearth  process,  893 
Ore  dressing,  900 
Organic  materials,  sources  of,  701 
Orthoclase,  807 
Orthophosphoric  acid.    See  Phosphoric 

acid,  ortho 
Osmic  acid,  838 
Osmium,  818,  838;  oxide,  838 
Osmosis,  711;  in  nature,  724 
Osmotic  pressure:,  measurement,  711; 

of  sugar  solution,  712;  theory  of,  713, 

7H 

Ostwald,  Wolfgang,  746 

Oxalic  acid,  665 

Oxidation,  325;  change  of  valence,  331, 
332,  334,  336;  intensity  of,  338 

Oxidation  and  reduction,  501-5;  cells, 
502;  electronic  explanation,  501;  by 
means  of  electric  current,  507; 
method  of  balancing  equations,  561; 
potentials,  506 

Oxides:  graphic  formula  of,  324;  of 
metallic  elements,  313;  of  non- 
metallic  elements,  313 

Oxyacetylene  torch,  315 

Oxygen:  constancy  of  concentration  in 
air,  766;  critical  temperature,  775; 


discovery,  305;  liquid,  309,  312,  777; 
minimum  weight,  63;  in  nature,  305; 
and  nitrogen,  544;  and  nitrogen, 
equilibrium,  566;  in  Periodic  Table, 
836;  from  plants,  311;  preparation 
of,  306-10;  properties  of,  16,  312,  313; 
uses  of,  315 

Oxone,  307 

Ozone,  316;  as  a  germicide,  317 

Paint:  drying  of,  364;  luminous,  859 

Paintings,  restoration  of,  347 

Palladium,  838 

Palmitic  acid,  678 

Palmitin,  679,  680 

Paraffine,  643;  series,  643 

Parkes  process,  912 

Pentathionic  acid,  628 

Perchlorates,  355 

Perchloric  acid,  355 

Perfume,  697 

Period  of  radioactive  substance,  864 

Periodic  law,  843,  844 

Periodic  system,  822-50 

Periodic  Table,  829;  anomalies  of,  848; 
arrangement  of  Harkins',  849;  history 
of,  844 

Permanganates,  343;  color  of,  in  solu- 
tion, 395 

Permonosulfuric  acid,  628 

Peroxides,  322;  graphic  formula  of,  324 

Perrin,  478,  703,  707 

Persulfuric  acid,  628 

Petroleum,  643 

Phenol,  672 

Phenolphthalein,  137,  440 

Phosgene,  695 

Phosphate  rock,  160,  580 

Phosphates:  precipitation  of,  452; 
qualitative  test  for,  597;  and  silver 
nitrate,  171, 597;  use  and  production, 

598 

Phosphine,  588 

Phosphonium  chloride,  588 

Phosphorescence,  584;  caused  by 
X-rays,  852 

Phosphoric  acid,  ortho:  and  albumin, 
597;  failure  to  precipitate  phos- 
phates, 452;  from  hydrolysis  of 
phosphorus  pentachloride,  247;  ioni- 
zation  of,  592;  normal  solution  of, 
595;  preparation,  158,  589,  590; 
properties,  159,  59°r92».  597 ;  and 
silver  nitrate,  597;  titration  of,  595. 
See  also  Metaphosphoric  acid;  Pyro- 
phosphoric  acid 

Phosphorous  acid,  588;  from  hydrolysis 
of  phosphorus  trichloride,  247 


644 


Introduction  to  General  Chemistry 


[References  are  to  sections,  not  pages] 


Phosphorus:  amorphous,  581;  and  bro- 
mine, 256;  and  chlorine,  247;  dis- 
covery of,  577;  effect  on  iron  and 
steel,  896;  and  iodine,  264;  manu- 
facture of,  580;  place  in  Periodic 
Table,  835;  red,  579,  581;  slowoxida- 
dation  of,  583;  white,  danger  in  use 
of,  585;  white,  physical  properties  of, 
578;  yellow,  578;  compounds,  588; 
pentabromide,  256;  pentachloride, 
247;  pentoxide,  247,  588;  sulfide, 
586;  tribromide,  256;  trichloride, 
247;  tri-iodide,  264;  tri-oxide,  588 

Photosynthesis,  691 

Phthallic  acid,  698 

Picric  acid,  673,  693 

Pig  boiling,  890 

Piperonal,  697 

Pitchblende,  874 

Plaster  of  Paris,  155 

Platinum:  adsorption  of  hydrogen  and 
oxygen,  731;  and  aqua  regia,  562; 
and  compounds,  818;  properties  and 
place  in  Periodic  Table,  838 

Polymer,  613 

Potash,  162;  caustic,  106;  monopoly  of 
Germany,  353.  See  also  Potassium 
hydroxide 

Potassium,  106;  action  on  water,  106; 
aluminum  sulfate,  175;  cupric 
chloride,  175;  flame,  784;  place  in 
Periodic  Table,  831,  848;  acid  tel- 
lurate,  815;  antimonyl  tartrate,  811; 
bicarbonate,  162;  bichromate  (see 
Dichromate);  bromide,  255;  use 
of,  260;  carbonate,  162;  chlorate: 
oxidizing  agent,  326;  preparation  of 
oxygen  from,  306;  preparation  and 
use,  352;  precipitation  of,  447; 
chromate,  346;  dichromate,  346;  and 
hydrogen  peroxide,  321;  and  hydro- 
gen sulfide,  346,  609;  and  sulfurous 
acid,  346;  ferricyanide,  665;  ferro- 
cyanide,  665;  hydrogen  tartrate,  665; 
hydroxide,  106;  iodide,  and  ferric 
chloride,  electronic  explanation  of, 
503;  nitrite,  from  potassium  nitrate, 
305;  perchlorate,  355;  preparation 
of  oxygen,  306;  permanganate  and 
ferrous  sulfate,  343;  and  hydrogen 
chloride,  235,  343;  and  hydrogen 
sulfide,  609;  oxidizing  agent,  326, 343; 
preparation  from  potassium  man- 
ganate,  837;  and  sulfurous  acid,  343; 
piatinic  chloride,  818;  tellurite,  815; 
uranyl  sulfate,  rays  from,  852 

Potato  starch,  preparation  of,  636 

Potential  energy,  368 


Potentials,  oxidation-reduction,  506 

Powder:  black,  571,  572;  smokeless, 
692 

Precipitates,  relative  stability  of  differ- 
ent forms,  762 

Precipitation,  154;  cause  of,  446;  classi- 
fication of,  452;  conditions  favoring, 
454;  of  electrolytes,  442;  by  sub- 
stance having  a  common  ion,  453 

Pressure:  critical,  775;  effect  of,  on 
gases,  4;  effect  on  system  in  equi- 
librium, 287;  partial,  113 

Priestley,  305 

Proof  spirit,  641 

Propane,  58,  643 

Proportions,  law  of  definite,  46 

Protecting  agents  for  colloids,  745 

Protein,  685 

Prout,  hypothesis,  851 

Prussic  acid,  665 

Ptyalin,  682 

Pure  substance,  23  0 

Pyrite,  6 10 

Pyrophoric  alloy,  820 

Pyrophosphoric  acid,  590 

Pyrosulfates,  624 

Pyroxaline,  694 

Quartz,  sand,  804 
Quartzite,  806 
Quicklime,  150 

Radical,  147;  acid,  377;  acid,  anions, 
391;  basic,  378;  phenyl,  670 

Radioactive  change,  theory  of,  865 

Radioactive  substances:  atomic  weights 
of,  869;  arrangement  of,  in  Periodic 
Table,  873;  valence  of,  872 

Radiothorium,  870 

Radium:  discovery  of,  855-56;  emana- 
tion, 863,  864;  heat  from,  862;  how 
measured  and  sold,  875;  origin  of, 
867;  place  in  Periodic  Table,  845, 
857;  ^  rays  of,  480;  series,  866; 
technical  production,  874;  total 
world-production,  874;  use  in  thera- 
peutics, 876 

Radium  clock,  860 

Rain,  formation  of,  772 

Ramsay,  791,  845 

Raoult,  718,  719 

Rare  earth  elements,  820;  place  in 
Periodic  Table,  846 

Rayleigh,  791 

Rays  of  radium,  480,  858-61.  See  also 
X-rays 

Reaction  velocity,  kinetic  hypothesis 
applied  to,  275-78 


Index 


645 


[References  are  to  sections,  not  pages] 


Reducing  agents,  intensity  of,  338 
Reduction,  327;  and  change  of  valence, 
335>  336;   electronic  explanation  of, 
501;  of  metals,  504;  of  non-metals, 

505 

Refractory  substance,  568 
Respiration,  314 
Reverb  eratory  furnace,  890 
Rhodium,  838 
Roasting,  of  ores,  902,  916 
Rock  crystal,  806 
Rock  salt,  24 
Roentgen  rays,  476 
Rose's  metal,  812 
Rowland,  468 
Rubber:  a  colloid,  760;  synthetic,  700; 

vulcanization  of,  604 
Rubidium,  823 
Ruby,  174 

Ruthenium,  place  in  Periodic  Table,  838 
Rutherford,  D.,  512 
Rutherford,  E.,  481,  865 

Safrole,  697 

Salicylic  acid,  697 

Sal  soda,  26 

Salt,  common:  composition,  45;  elec- 
trolysis of,  238;  preparation,  41; 
purification  from  rock  salt,  24 

Salt,  parts  of,  376 

Saltpeter,  Chile,  104;  occurrence,  540 

Salts:  acid,  102;  active  parts  of,  376; 
a  class  of  chemical  substances,  92; 
hydrolysis  of,  436;  ionization  of, 
410;  little  soluble  of  strong  acids, 
not  dissolved  by  acids,  458;  neutral, 
102;  primary,  secondary,  and  terti- 
ary, 159 

Sandpaper,  806 

Sandstone,  806 

Saponification,  679 

Sapphire,  174 

Sassafras,  oil  of,  697 

Scandium,  discovery  of,  845 

Scheele,  305 

Scurvy,  689 

Seaweeds,  iodine  from,  261 

Selenium,  and  its  compounds,  815,  836 

Series:  displacement  of  metals,  492; 
of  non-metals,  488 

Siedentopf,  706 

Siemens-Martin  process,  893 

Silica,  806 

Silicates,  807 

Silicic  acid,  807;  a  gel,  761 

Silico-chloroform,  805 

Silico-ethane,  805 

Silico-mesoxalic  acid,  805 


Sili co-methane,  805 

Silico-oxalic  acid,  805 

Silicon:  compared  to  carbon  com- 
pounds, 805 ;  and  compounds,  804-8, 
827;  place  in  Periodic  Table,  834; 
dioxide,  806;  fluoride,  270;  hexa- 
chloride,  805;  tetrachloride,  805 

Silk,  artificial,  694 

Silver,  168;  colloidal,  744;  metallurgy 
of,  926;  place  in  Periodic  Table,  831; 
production  of,  881;  test  for,  169; 
acetate:  precipitation,  452;  solution 
by  nitric  acid,  456;  ammonium  ion, 
533-34;  effect  of  acids  on,  536; 
ammonium  salts,  535,  537;  azid, 
531;  bromide,  257;  use,  260;  chloride, 
169;  solubility  in  ammonia  solution, 
532-35;  hydroxide,  172;  nitrate,  1 68; 
electrolysis  of  ^ solution,  387;  and 
hydrochloric  acid  in  the  presence  of 
gelatine,  745 ;  oxide,  172;  phosphate, 
171,  597;  salts,  reaction  with  chloride, 
382;  sulfate,  170;  sulfide,  precipita- 
tion of,  452;  thiosulfate,  complex 
ion,  626 

Slag,  885,  888,  903,  911 

Smithsonite,  915 

Smoke  screens,  587 

Soap:  cleansing  action,  751-54;  com- 
position, 162,  678 

Soda,  26;  caustic,  and  hydrochloric 
acid,  89;  caustic,  from  sodium  and 
water,  40;  washing,  26 

Soda  water,  633 

Sodamide,  527 

Soddy,  481,  865 

Sodium:  and  chlorine,  theory  of  union 
of,  485;  flame,  784;  place  in  Periodic 
Table,  831;  spectrum  of,  786;  and 
water,  40;  acetate,  157;  action  of 
water  on,  435;  and  hydrochloric 
acid  (ionic  theory),  424;  aluminate, 
177;  aluminum  fluoride,  267;  alu- 
minum sulfate,  175;  antimony  flu- 
oride, 811;  benzoate,  671;  bicar- 
bonate, 161;  bromide,  255;  use  of, 
260;  carbonate,  161;  chlorate,  353; 
chloride,  action  on  sulfuric  acid,  460; 
electrolysis  of  solution,  385;  pre- 
cipitated by  hydrochloric  ^acid,  453; 
fluoride,  269,  270;  fluosilicate,  272; 
hydrogen  sulfate,  98,  100;  action  of 
hydrochloric  acid  on,  289;  hydro- 
sulfite,  628;  hydroxide,  88;  and 
acetic  acid  (ionic  theory),  434;  and 
ammonium  chloride,  426,  462;  and 
hydrochloric  acid,  421;  hypochlorite, 
350;  metaphosphate,  594;  nitrate^ 


646 


Introduction  to  General  Chemistry 


[References  are  to  sections,  not  pages] 


104;  decomposition  by  heat,  565; 
nitrite,  513;  oleate,  work  of  Harkins 
on,  754;  perchlorate,  355;  peroxide, 
307;  phosphates,  159,  594;  pyrophos- 
phate,  594;  silicate,  270,  807;  silver 
thiosulfate,  626;  stannite,  809;  sul- 
fate,  94;  hydrate  of,  94,  96,  97; 
sulfides,  607;  tetraborate,  general 
properties,  801,  803;  tetrathionate, 
625;  thiosulfate,  625;  tungstate,  814; 
zincate,  177 

Solder,  809,  913 

Solid  state,  theory  of,  203 

Solubility:  curves,  132;  effect  of  crys- 
talline form  on,  133;  effect  of  tem- 
perature on,  132-34;  gases  in  liquids, 
125;  and  heat  of  solution,  134; 
liquids  in  liquids,  124;  of  molecules, 
445;  of  solids,  122;  two  kinds,  135 

Solution,  120;  compared  to  gas,  708; 
comparison  with  suspension,  708; 
concentration  of,  121;  electrolysis 
of,  385;  heat  of,  127;  kinetic  theory 
of,  443;  of  little  soluble  electrolytes, 
455;  normal,  136-38,  595;  saturated, 
12-2;  supersaturated,  123;  van'tHoff's 
theory,  715 

Soot,  formation  in  name,  782 

Spark  spectra,  786 

Specific  gravity,  no 

Specific  heat,  in;  and  symbol  weight, 
229 

Spectra:  bright-line,  786;  continuous, 
793;  dark-line,  789;  of  gases  and 
volatile  substances,  787 

Spectroscope,  785 

Spectroscopic  analysis,  786 

Spelter,  148,915 

Sphalerite,  915 

Spinthariscope,  859 

Spirit,  proof,  641 

Stannic  compounds,  809 

Stannous  compounds,  809 

Stannum,  809 

Starch:  glucose  from,  636;  occurrence 
of,  636;  properties  of,  637,  638;  test 
for,  637 

Steam,  8;  action  on  iron  (chemical 
equilibrium),  290;  burns  by,  117; 
production,  fuel  for,  360;  use  for 
heating,  116 

Stearic  acid,  678 

Stearin,  680 

Steel,  8835.;  alloys,  898;  Bessemer, 
891;  crucible,  894;  electric,  894; 
high  speed,  898;  stainless,  898; 
tempering  of,  897 


Stibine,  8n 

Stoney,  466 

Storage  battery,  914 

Strontium,  832;  flame,  784 

Strutt,  860 

Sublimation,  179 

Subscripts,  chemical,  61 

Substance,  22,  23;  properties  of  a,  108 

Sugars,  structure  of,  683 

Sulfates,  624;  test  for,  164,  380 

Sulfides:  precipitation^  of,  608;  pre- 
cipitation by  ammonium  sulfide,  452 

Sulfites,  612 

Sulfocyanates,  665 

Sulfur:  {commercial  importance  of,  604; 
consumption  of,  604;  effect  on  iron 
and  steel,  896;  flowers  of,  603;  heat 
of  combustion,  ^35  7;  monoclinic,  600; 
place  in  Periodic  Table,  836;  plastic, 
600;  properties  of,  600,  60 1 ;  rhombic, 
600;  rock,  603;  roll,  603;  sources  of, 
602;  dioxide:  critical  temperature, 
775;  preparation  and  properties,  610; 
and  water,  equilibrium,  286;  mono- 
chloride,  600;  trioxide,  613. 

Sulfuric  acid:  and  aluminum  hydroxide, 
175;  anhydride  of,  613;  and  copper 
oxide,  165;  dilute,  action  on  iron, 
173;  dilute,  action  on  magnesium, 
149;  dilute,  action  on  zinc,  149;  fum- 
ing, 613;  importance  of,  614;  and 
magnesium  hydroxide,  144;  and 
magnesium  oxide,  145;  manufacture 
of,  615;  neutralization  of,  94;  prop- 
erties, 92,  ^620-23;  and  sodium 
chloride  (ionic  theory),  460 

Sulfurous  acid,  340,  6n;  equilibrium, 
sulfur  dioxide  and  water,  286;  and 
iodic  acid,  277 

Sun:  composition  of,  790;  source  of 
energy,  374 

Supercooling,  206 

Superphosphate,  160 

Surface,  increase  with  subdivision,  727 

Suspensions,  i2oj  laws  governing,  707 

Suspensoids,  735,  747;   of  commercial 
importance,   750;    influence  of   the 
charge  on  stability  of,  740 
Symbol  weights:  accuracy  of,  222;  how 
found,  224,  228;   oxygen  basis,  223; 
product  of,  and  specific  heat,  229 
Symbols:  chemical,  61;  and  minimum 

weights,  66 
Synthesis,  31;  organic,  696,  698 

Table  sugar,  682 
Talc,  807 


Index 


647 


[References  are  to  sections,  not  pages] 


Tannin,  used  in  making  red  gold  solu- 
tions, 746 
Tantalum,  835 
Tantiron,  804 
Tartar  emetic,  811 
Tartaric  acid,  665 
Tellurium  and  its  compounds,  815,  836, 

848 

Temperature:  absolute,  6;  critical, 
775;  effect  of  change  of,  on  equi- 
librium, 288;  effect  of  change  of,  on 
gases,  5;  effect  of  change  of,  on  solu- 
bility, 134,  288 

Tetramethyl  silico-methane,  805 
Tetrathionic  acid,  628 
Thermal  Unit,  British,  358 
Thermite,  330,  925 
Thermometer:  centigrade,  6;  gas,  6 
Thermos  bottles,  777 
Thiocyanate  (sulfocyanate),  665;    and 

ferric  chloride,  280,  415 
Thomas-Gilchrist  process,  892 
Thorium:  and  its  active  products,  870; 
and  its1  compounds,  820;  in  monazite, 
820;    place  in  Periodic  Table,  834; 
properties,  etc.,  834  _ 
Thorium   oxide,   use   in  gas   mantles, 

301 

Thorium  X,  870 

Tin:    and   its   compounds,    809,    834; 
amalgam,    foil,    plate,    ware,    809; 
place  in  Periodic  Table,  834 
Titanium,  834 

Titration,  137;  indicators  for,  440 
T.N.T.,  673,  693 
Transition  point,  600 
Transmutation  of  elements,  851 
Trimethyl  amine,  59,  658 
Tri  nitro  phenol,  673,  693 
Tri  nitro  toluene,  673,  693 
Tripoli,  806 
Trithionic  acid,  628 

Tungsten:  and  its  compounds,  814, 
836;  use  in  incandescent  filaments, 
814 

Turkey  red,  698 
Turpentine,  248 
Tuyere,  886 
Tw.    SeeTwadell 

Twadell,  scale  of  specific  gravity,  618 
Tyndall,  J.,  748;  effect,  748 
Type  metal,  811 

Ultramicro scope,  703,  706 
Unit  volume,  the  chemical,  65 
Univalent,  146 
Unsaturated  compounds,  527,  663 


Uranium:      parent    of    radium,     867; 

properties     and     compounds,     836; 

series,  868 
Uranyl  nitrate,  836 
Urea,  659;  synthesis  of,  696 

Valence,  146;  atoms  and  radicals,  183; 
change  of,  in  oxidation  and  reduction, 
33I-36;  electrical  nature  of,  484, 
880;  and  Periodic  Table,  850;  posi- 
tive and  negative,  840;  of  radicals, 
147;  and  the  structure  of  inorganic 
molecules,  839;  table  of,  183 

Valeric  acid,  680 

Vanadium,  835 

Vanilla,  697 

Vanillin,  697 

Van't  Hoff,  715,  719 

Vapor,  equilibrium  with  liquid,  201 

Vapor  pressure:  correction  of  volume 
of  gas  for,  114;  of  hydrates,  131; 
liquids  and  solids,  114;  lowering  of, 
by  dissolved  substance,  129;  theory 
of,  200;  of  water,  112 

Varnish,  drying  of,  364 

Vaseline,  643 

Velocity  of  reactions,  275 

Vinegar,  157;  cider,  653;  plant,  653 

Vitamines,  689 

Vitriol:  blue,  497;  oil  of,  619 

Volt,  499 

Volume,  atomic,  823 

Warfare,  materials  for,  chemical,  695 

Water:  adsorption  by  glass,  730;  com- 
position of,  36;  distilled,  23;  elec- 
trolysis of,  27,  295;  formation  of, 
139;  formula  of,  70;  hard,  156,  803; 
ionization  of,  420;  a  product  of 
combustion,  20;  properties,  108-12; 
reactions  of,  140;  and  sodium,  40; 
synthesis  of,  37;  vapor,  in  air,  766; 
vapor  pressure,  112 

Water  glass,  808 

Water  soluble  B,  689 

Weight,  equivalent,  403 

Weights:  atomic,  800  (see  of  so  inside 
of  back  cover) ;  law  of  minimum  and 
multiple,  64,  212;  minimum  of  ele- 
ments, 63;  minimum  and  symbols, 
66;  symbol  and  formula  weights,  74 

Whetstone,  806 

White  lead,  913 

Willemite,  807, 9 1 5 ;  phosphorescent,  86 1 

Winkler,  845 

Wintergreen,  oil  of,  697 

WBhler,  696 

Wollaston,  790 


648 


Introduction  to  General  Chemistry 


[References  are  to  sections,  not  pages] 


Wollastonite,  807  '  *• J»  j l 

Wood  alcohol,  645 
Wood's  metal,  812 
Work,  368 f;r  •         •    ,   , 
Wolframite;  8i} 

Xenon:  discovery  and  properties,  798; 

and  family/825 
X-rays,  476;  and  crystal  structure,  204, 

879;  spectra,  877 
Xylene,  668 

Yeast,  640 


Zinc:  and  concentrated  sulfuric  acid, 
622;  and  dilute  sulfuric  acid,  149; 
electrolytic,  918;  and  hydrochloric 
acid,  149;  metallurgy  of,  916-18; 
and  nitric  acid,  561;  place  in  Periodic 
Table,  832;  production,  881,  915; 
properties  and  use,  148,  919;  chlo- 
ride, 148;  nitrate,  148;  oxide,  148, 
919;  sulfate,  148;  sulfide,  919;  and 
alpha  rays,  859;  precipitation  from 
solution,  452 

Zirconium,  and  compounds,  834 

Zsigmondy,  706 


•        -. 


PERIODIC  TABLE  ACCORDING  TO  ATOMIC  NUMBERS,  H:i 


Pe- 
riod 

0 

A         B 

II 
A          B 

III 
A          B 

IV 
A          B 

V 
A          B 

VI 
A         B 

VII 
A          B 

VIII 

I 

2 

He 

Li3 

Be4 

5 
B 

6 

c 

N 

8 
O 

9 

i 

2 

10 

Ne 

IO 

Na 

12 

Mg 

I3, 

14 
Si 

15 
p 

16 

g 

17 
Cl 

18 
A 

K19 

20 

Ca 

21 

Sc 

Ti22 

v23 

Cr24 

Mn5 

26    27    28 

Fe    Co    Ni 

3 

29Cu 

3°Zn 

3IGa 

32Ge 

33 

As 

34Se 

36 
Kr  < 

37 
Rb 

s,38 

y» 

Z/° 

3 

Cb 

42 

Mo 

?43 

44     45     46 
Ru    Rh   Pd 

4 

47 
Ag 

48cd 

49 

In 

5°Sn 

51  sb 

52 
Te 

53 

& 

55 
Cs 

Ba56 

£1 

£ 

f? 

60 
Nd 

61 
? 

5 

62 
Sa 

63 

Eu 

64 
Gd 

65 

Tb 

66 
Dy 

67 
Ho 

67 
Er 

JL 

70 
? 

7i 
Yb 

11 

Ta" 

74 
W 

75 

76     77     78 
Os    Ir    Pt 

79Au 

80 
Hg 

81 
Tl 

82 
Pb 

83  Bi 

84 

85 

7 

86 

Nt 

?  8? 

88 
Ra 

?  8Q 

90 
Th 

91 

? 

u92 

THIS  BOOK  IS  DUE  ON  THE  LAST  DATE 
STAMPED  BELOW 

AN  INITIAL  FINE  OF  25  CENTS 

WILL  BE  ASSESSED  FOR  FAILURE  TO  RETURN 
THIS  BOOK  ON  THE  DATE  DUE.  THE  PENALTY 
WILL  INCREASE  TO  SO  CENTS  ON  THE  FOURTH 
DAY  AND  TO  $1.OO  ON  THE  SEVENTH  DAY 
OVERDUE. 


!)EC    18  1933 


. 

/? 


or 


IB  28  1937 


ffM  *  nn. 

5 

iff    -\    W3<, 

Af(    I--    »K5 

r,    • 

• 

ffA*f     !  j  & 

frr  3   r 


FEB     3    1938 


AUi 


MAR  1JJ  1940 


^IS^^i" 


:   ^ggsao^/ 

^5Ss^ 

Berkeley 


'™i'y°fc£Kraill 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 


